Objectives

When you have completed this chapter you should be able to:

• Define the differences between elements, compounds and mixtures
• Explain the difference between organic and inorganic chemistry
• Describe the structure of an atom
• Show how the periodic table is formed and describe the different properties of elements
• Show how ions are formed and understand both covalent and ionic bonding
• Understand chemical formulae and the use of chemical equations
• Explain atomic weight and molar concentrations
• Discuss the difference between acids and bases and explain oxidation-reduction

1.1 Introduction

1.1.1 Elements, compounds and mixtures

All substances may be divided into three classifications: elements, compounds and mixtures.

The simplest forms of matter are elements — substances that cannot be split into simpler substances by a chemical reaction. As shown in Figure 1.1, there are over 100 known elements classified in the periodic table and these form the building blocks of all chemicals.

A compound is formed by two or more elements which are always combined in the same fixed ratio. Thus, for example, water is a compound of hydrogen and oxygen, shown by its chemical formula H2O, which is always formed from two parts hydrogen and one part oxygen (Figure 1.2). Compounds are often difficult to split into their elements and can only be separated by chemical reactions.

The smallest particle of an element or compound that normally exists by itself and still retains its properties is called a molecule. Normally a molecule consists of two or more atoms — sometimes even thousands. However, as we shall see later, a molecule can also exist as a single atom called a monatomic molecule.

Elements and compounds are pure substances whose composition is always the same. In reality, pure substances are rarely encountered and most are mixtures of compounds or elements that are not chemically combined and in which the proportions of each element or compound are not fixed (Figure 1.3).

Since the elements or compounds making up the mixture each keep their own properties they can usually be separated fairly easily by physical means.

When a mixture has the same properties and composition throughout, (e.g. a mixture of sugar and water in which the sugar is thoroughly dissolved in the water) it is called a homogeneous mixture or a solution.

When the composition and properties of a mixture vary from one part to the other it is called heterogeneous. This might take the form of a suspension in which fine particles of a solid are suspended in a liquid and are not dissolved. Another type of heterogeneous mixture is a two-phase mixture: e.g. a mixture of oil and water in which the oil floats on the water as a separate layer.

1.1.2 Organic and inorganic chemistry

In the study of elements and their components, a distinction is made between what is termed organic and inorganic chemistry.

Originally organic chemistry concentrated on the study of substances found in living organisms. Now, however, it extends to cover all (with a few exceptions) of the compounds of carbon. The vast number of organic compounds (well over two million) include plastics, fibres, drugs, cosmetics, insecticides, foods, etc., and are made possible by the ability of the carbon atom to bond to itself and form a virtually unlimited chain or ring.

Inorganic chemistry is the study of that which is left over — those substances that do not form bonds with carbon.

1.2 Atomic structure

At one time the atom was considered to be the ‘fundamental particle’ — the smallest bit of matter that could be conceived. And indeed, since all atoms of any given element behave in the same way chemically, the atom is the smallest entity to be considered from a chemical viewpoint.

In considering the basic construction of an atom it is usual to consider it as being composed of three elementary sub-atomic particles: protons, neutrons and electrons. It is the number of protons, neutrons and electrons that any atom contains that distinguishes one element from another.

The main mass of the atom comprises the nucleus, which consists of protons and neutrons. The proton has a positive charge and a relative mass of 1 whilst the neutron has a similar mass but no charge.

The nucleus is surrounded by a cloud of electrons each having a negative charge and a mass nearly 2000 times smaller than the proton. In its normal state, an atom has the same number of electrons as protons and is, therefore, electrically neutral.

The simplest structure is the hydrogen atom, which comprises a single proton and a single electron (Figure 1.4).

An element is determined by the number of protons in the nucleus — its atomic number. Since the atomic number defines the element, any atom with 6 protons is carbon, irrespective of the number of neutron or electrons (normally 6 of each).

However, because the number of neutrons in the nucleus may not always be the same as the number of protons, an atom is also defined in terms of the total number of protons and neutrons in the nucleus — called its mass number (not to be confused with its atomic mass). Carbon, for example would normally have 6 protons (atomic number = 6) and 6 neutrons giving a mass number of 12. However, carbon can also exist in other forms, called isotopes, in which the nucleus contains 7 or even 8 neutrons — giving a mass number of 13 or 14 respectively.

The three carbon isotopes are distinguished from one another by writing the mass number after the name of symbol of the element. Thus, as shown in Figure 1.5, carbon with 6 neutrons is called Carbon-12; with 7 neutrons is called Carbon-13; and with 8 neutrons is called Carbon-14.

Electron shells

With the electrons forming a cloud around the nucleus, the position of any electron is defined in terms of the probability of finding it at some distance from the nucleus. Consequently, it is convenient to visualise the electrons moving around the nucleus of an atom in the same way as planets move around the sun.

This solar-system model visualises the electrons arranged in definite shells or energy levels. Each shell has an upper limit to the number of electrons that it can accommodate and each is built up in a regular fashion from the first shell to a total of seven shells. The first shell can accommodate 1 or 2 electrons; the second shell a maximum of 8; the third a maximum of 18; the fourth 32 and so on with successive shells holding still larger numbers.

In the study of chemistry, it is the electrons in the outermost shell — those which are added last to the atom’s structure — that determine the chemical behavior of the atom. This outer shell is called the valance shell and the electrons in it are called valance electrons. For the first two shells (the first 10 elements) matters are simple and each shell is filled before a fresh one is started. However, in later elements the outer shell develops before the previous one is completed.

Chemical symbols

As shown in Figure 1.1, each element has its own unique symbol that, in most cases, is formed from one or two letters of the element’s English name: e.g. H for hydrogen; He for helium; Li for lithium. Note that the first letter is always capitalised and, where there is a second letter, this is always written in lower case. The symbols of some elements are derived from their Latin names (Table 1.1): Na for sodium from its Latin name Natrium; Au for gold from its Latin name Aurum; Pb for lead from its Latin name Plumbum; etc. An exception to this derivation is tungsten whose symbol W is derived from its German name Wolfram.

1.3 Periodic table

The periodic table (Figure 1.1) is an arrangement of the elements, in order of their increasing atomic numbers, in which a wide range of chemical and physical information is arranged in a systematic way.

As shown, there are seven rows, called periods, with the atomic number increasing by one element from left to right.

The vertical columns or groups are numbered using roman numerals and separate the elements into families having the same number of electrons in the outer shell and similar chemical properties.

As shown in Figure 1.7, about 75% of the elements are made up of what are called metals. Except for mercury, metals are solid and are characterised by high conductivity (both electrical and thermal); high malleability (the ability to be formed or hammered into different shapes); and high ductility (the ability to be drawn into a thin wire).

Moving to the right, the elements gradually become less metallic and, on the extreme right (Figure 1.8), are termed non-metals. Non-metals can be solid (carbon being the most commonly observed) liquid (bromine) or gas (oxygen, nitrogen) and generally have poor conductivity, malleability and ductility.

Somewhere in between the metallic and non-metallic elements are the metalloids (Figure 1.9) that behave mainly like non-metals except that their electrical conductivity, although not good, is closer to metals. These elements are thus often termed semiconductors.

1.4 Properties of elements

Elements to the extreme left of the periodic table (Figure 1.10), in Group I are, with the exception of hydrogen*, called alkali metals — metals that react with water to form alkaline solutions. All have only one valence electron (one electron in their outer shell) which is easily lost in reactions. Indeed, moving down the group from lithium, sodium, potassium, rubidium, caesium, and francium, their reaction with water becomes increasingly more violent and all need to be stored under oil because of this reactivity.

*Hydrogen owes its location in Group I to its electron configuration rather than its chemical properties.

The metals of Group II (Figure 1.11) have two fairly loosely bound valance electrons and, although they are also strongly reactive, are not as reactive as Group I elements. Group II elements tend to be found as naturally occurring mineral deposits in the ground or in the sea and are, therefore called the alkaline earth metals.

At the far right hand side of the periodic table (Figure 1.12) are the Group VIII (also know as Group O) elements that form the noble or inert gases. The term noble indicates that the element is chemically inert or inactive.

Noble gases – helium, neon, argon, krypton, xenon and radon – all have a full outer shell which accounts for their extreme stability and unwillingness to form compounds. Indeed, at normal temperatures, the lighter noble gases form no compounds whatsoever and exist as monatomic (single-atom) molecules.

In contrast, gases such as hydrogen, nitrogen, oxygen, fluorine, and chlorine react with each other to form diatomic (two-atom) molecules (Figure 1.13) e.g. H2, N2, O2, F2, and Cl2. The Non-metallic elements bromine and iodine also exist as diatomic molecules: Br2 and I2.

1.5 Formation of ions

Earlier we stated that the nature of the element is determined by its atomic number – the number of protons within the nucleus – and that it can gain or lose electrons. We also saw how the electrons in an atom are arranged in shells.

Because of the increasing distance of successive shells from the nucleus, the outer electrons are further away from the nucleus and are thus held less tightly. The result is that less energy is required to remove an electron from the 2nd shell than from the 1st shell and still less to remove an electron from the 3rd shell than the 2nd shell. Thus, referring to the periodic table, the energy required to remove an electron from the outer shell of an atom decreases from the top to the bottom (Figure 1.15).

In addition, elements to the left of the table have loosely held valence electrons that are easily lost. Now, moving from left to right, as electrons are added to the outer shell, an effectively increasing charge binds them more and more tightly to the nucleus.

It can thus be seen that it is very difficult for elements in the upper right-hand corner to lose electrons and indeed elements with nearly completed shells will tend to attract electrons.

Figure 1.16 shows a sodium atom (atomic number 11) in which the first and second shells are complete and there is one valence electron. Assume now that the atom loses the electron in its outmost shell. What will be the result?

Firstly, since the nucleus has been unaffected and it still has 11 protons (as well as 11 neutrons) it will continue to remain sodium.

The main effect will be that the total charge of the atom will no longer be neutral but will be unbalanced by the absence of the negatively charged electron. The result is that the atom will now have a net positive charge. Since, by definition, an atom is neutral, the change in its charge by losing the electron has, in effect created a new particle called an ion. A positively charged ion is called a cation (pronounced CAT-ion).

The loss of the electron (e^–) from the sodium atom creates an ion with a net charge of +1 and the particle’s change in status from a sodium atom to a sodium cation (Figure 1.17) could thus be indicated by the symbol Na^1+. In practice the 1 is implied and the symbol is thus written Na⁺.

If a magnesium atom (atomic number 12) with two valence electrons (Figure 1.16) should lose them, then the resulting cation will be unbalanced by the absence of two electrons and thus have a net positive charge of +2. In this case the particle’s change in status from a magnesium atom to a magnesium cation (Figure 1.18) is indicated by the symbol Mg ²⁺.

If we examine the structure of a chlorine atom (atomic number = 17), as depicted in Figure 1.20, it can be seen that the outer (third) shell contains only 7 electrons. And if it should capture an electron and make its outer shell complete (Figure 1.21) its overall charge will become negative. A negatively charged ion is called an anion (pronounced AN-ion).

The gain of the electron (e^– ) into the outer shell of the chlorine atom, creates an ion with a net charge of -1 and the particle’s new status as a chlorine anion is indicated by the symbol Cl^– — again with the 1 being implied.

1.6 Bonding

Earlier, we described a compound as a combination of two or more elements. The elements that are combined to form compounds are held together by bonding forces called chemical bonds. The two major bonding forces are ionic bonding and covalent bonding.

1.6.1 Ionic bonding

Ionic bonding is based on the fact that opposite charges attract!

We have already seen how Group I and II elements have loosely held valence electrons that are easily lost and how elements with nearly completed shells will tend to attract electrons.

Thus a sodium atom with a single valence electron will give it up easily to, for example, a chlorine atom with 7 valence electrons (Figure 1.22).

The sodium cation, with a net charge of +1 and a chlorine anion having a net charge of -1 are attracted to each other and thus form an ionic bond. This ionic bond, formed between a sodium cation (Na⁺) and a chlorine anion (Cl^–), results in the compound sodium chloride (Na⁺Cl^–) or common salt (Figure 1.23). In practice the ionic charges are omitted from the formula which then becomes NaCl.

Another example of an ionic compound formed in this manner is that of magnesium and oxygen. Here, a magnesium cation (Mg²⁺) having a net charge of +2, combines with an oxygen anion (O^2-) having a net charge of -2. The ionic bonding results in the compound magnesium oxide (Mg²⁺O^2-) — normally written as MgO.

Ionic compounds thus exist as a collection of electrically charged positive and negative ions – with each positive ion surrounding itself with as many negative ions as it can, and each negative ion surrounding itself with as many positive ions as it can.

The resulting structure consists of a regular pattern of alternate positive and negative ions, in which no individual molecules can be identified. This regular arrangement of positive and negative ions continues indefinitely in three dimensions throughout the whole structure and is stacked together to form a crystal lattice (Figure 1.24).

Such ionic crystals are hard and brittle with high melting points. Because, at room temperature, the ions are bound and are not free to move, ionic compounds do not normally conduct electricity. However, when melted or dissolved in water, the ions are freed and the materials become conductive.

The ionic compounds formed in this fashion between metals and non-metals are called salts.

1.6.2 Covalent bonding

Covalent bonding is a more complex arrangement than ionic bonding. Here, the valence electrons are actually shared between the atoms so that each acquires a stable outer shell.

When two atoms of hydrogen are brought together their electrons begin to feel the attraction of both nuclei (Figure 1.25). As the distance continues to decease, there is an increasing probability of finding either electron near either nucleus and at some point, each hydrogen atom shares an electron and the two are thus bonded together to form an H₂ molecule. A less accurate, but more easily visualised, analogy is that the electrons form a figure-of-eight path (Figure 1.26) around the two nuclei to form the H₂ molecule.

In ionic bonding the ions exist separately. In covalent bonding, however the net molecular charge is neutral. In the case of hydrogen, the covalent bond is formed by a single pair of electrons – with the pair usually connected by a single dash (-):

H – H

Another example of a single-pair covalent bond is chlorine (Figure 1.27). When chlorine atoms approach each other (each with an electron valence of seven) covalent bonding occurs when a pair of electrons are shared (one from each atom) so that each atom now has a stable outer shell of 8 electrons. Again the molecule formed by this single-pair bond is represented by: Cl – Cl

Yet another single-pair covalent bond can be formed by, for example, hydrogen and chlorine (Figure 1.28). Here, the pair is made up by one electron from each atom to form a stable outer shell for each. The resulting compound, hydrogen chloride (HCl) is depicted by: H – Cl

More than one pair of electrons can form a covalent bond. An oxygen molecule, for example, (with six valence electrons) is formed when two pairs of electrons are shared – again resulting in a stable outer shell of eight electrons for each atom. Here, the double pair is represented by a double dash (=):

O = O

In the case of nitrogen, with five valence electrons, a molecule is formed by a triple bond in which three pairs of electrons (three electrons from each atom) give rise to a stable outer shell for each atom. Here, the triple pair is represented by a triple dash (=):

N = N

Chemical formulae

The most commonly used chemical formula is the molecular formula, which gives the actual composition of a molecule. The molecular formula H₂O, for example, indicates that a water molecule comprises two H atoms and one O atom. Similarly, the formula for ethene is C₂H₄ — indicating that an ethene molecule is made up of two carbon atoms and four hydrogen atoms.

At this point a distinction needs to be drawn between the molecular formula, stated above, and what is termed an empirical formula.

Whilst the molecular formula gives the composition of a molecule, the empirical formula expresses the composition of a substance in terms of whole-number ratios.

Figure 1.29 shows the structure of an ethene molecule. The empirical formula would be written CH₂ — indicating the ratio of atoms i.e. there are two hydrogen atoms for each carbon atom. The molecular formula, on the other hand (written C₂H₄), indicates that a molecule of ethene comprises two carbon atoms and four hydrogen atoms.

For even more information the shortened structural formula is used. This shows the sequence of groups or atoms in a molecule. And for more information still, the full structural or displayed formula shows all the bonds (both single and double) that make up the molecule.

It should be noted that in organic chemistry, even more information than that provided by the displayed formula is required. For example, isomers are compounds having the same molecular formula but different arrangements of atoms in their molecules.

Often, such compounds can only be distinguished from each other using stereochemical formulae that show the structure of the molecule in 3-D.

Thus, whilst the molecular formula for butene is C₄H₈ it has two different stereochemical formulae (Figure 1.30).

We have already seen how some elements such as oxygen and nitrogen occur as simple molecules even when not combined with other elements. These give the formula O₂ and N₂ respectively.

1.7 Chemical equations

A chemical equation gives a ‘before and after’ picture of a chemical reaction.

In its most basic form, it may be written out:

Sodium + Water → Sodium hydroxide + Hydrogen

The two substances that produce the reaction, and are called reactants, are written on the left hand side of the equation and the result or products of the reaction are written on the right hand side. The + sign stands for ‘react with’ and the arrow (→) stands for ‘reacts to yield’. Thus, the above equation indicates that sodium reacts with water to yield sodium hydroxide and hydrogen.

In practice, the names of the materials involved in the reaction are not written out in full but are replaced by their chemical formulae:

Na + H₂O → NaOH + H₂

Whilst this equation shows the reaction that takes place, it does not convey the complete picture since closer examination shows that it is chemically unbalanced — i.e. the number of atoms on the one side is not the same as the other (Table 1.2).

In order to balance the equation it may be rewritten as shown with the number in front of each equation, called the coefficient, indicating the number of atoms :

2Na + 2H₂O → 2NaOH + H₂

Now, an audit, as shown in Table 1.3, show that there is a balance.

The equation may be extended even further by including what are termed state symbols after each formula that indicate the physical state of each molecule. The state symbols (written in brackets) are:

(s)	=	solid
(l)	=	liquid
(aq)	=	aqueous (solution in water)
(g)	=	gas

Thus, the sodium reaction with water would be:

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂ (g)

Finally, when studying electrolytes in solution, the chemical properties are best represented by an ionic equation which is only concerned with the ions of the participant substances in the reaction.

Thus the reaction:

NaOH(aq) + HCl(aq) → NaCl(aq) + H₂O(l)

may be rewritten:

Na⁺OH^–(aq) + H⁺Cl^–(aq) → Na⁺Cl^–(aq) + H₂O(l)

Here, Na⁺, OH^– , Cl^– and H⁺ are all ions. Since Na⁺ and Cl^– appear on both sides of the equation, they are uninvolved the ionic reaction and are thus omitted. Now the ionic equation is simply:

Na⁺OH^–(aq) + H⁺Cl^–(aq) → Na⁺Cl^–(aq) + H₂O(l)
OH^–(aq) + H⁺ (aq) → H₂O(l)

1.8 Naming compounds

Many compounds have trivial names —names that give little information about the structure of the compound. These include names such as salt (NaCl), borax (Na₂B₄O₇·10H₂O), chalk (CaCO₃), or even proprietary names such as Teflon (F(CF₂)_nF). In addition, there are other chemical compounds whose names are universally associated with a chemical formula (e.g. water (H₂O)). However, most of these names are often difficult to interpret. As a result, most compounds are named using a systematic internationally agreed system.

The compounds of metals and non-metals containing, for example, two different elements are named by first taking the name of the metallic element followed by the main part of the name of the non-metal, which is then modified with the suffix -ide. Thus, the compound of sodium and chlorine takes the metallic element first (sodium) followed by the modified form of chlorine (chloride): i.e. sodium chloride (NaCl).

Other such compounds include:

CaS — calcium sulfide;

MgO — magnesium oxide;

SiN — silicon nitride; and

ZnS — zinc sulfide

When more than two atoms are involved use can be made of a prefix:

mono-	1
di-	2
tri-	3
tetra-	4
penta-	5
hexa-	6
hepta-	7
octa-	8
nona-	9
deca-	10

Thus we get:

CO — carbon monoxide;

CO₂ — carbon dioxide

NO₂ —nitrogen dioxide;

CS₂ — carbon disulfide;

SF₆ — sulfur hexafluoride;

GeCl₄ — germanium tetrachloride; and

N₂O₄ — dinitrogen tetraoxide.

Another frequently used suffix is -ate which usually indicates the presence of oxygen. Thus: nitrate, NO^3-; sulfate, SO₄^2-; and phosphate, PO₄^3- .

The suffix -ite indicates fewer oxygen atoms than in the corresponding -ate ion, with the prefix hypo- used with the suffix -ite indicating still fewer. The prefix per- indicates more oxygen, or less negative charge, than the corresponding -ate ion. Thus:

Chlorate	ClO3–
Chlorite	ClO2–
Hypochlorite		ClO–
Perchlorate		ClO4–

Certain metals can be found in more than one ionic state. Iron, for example, occurs as either Fe²⁺ or Fe³⁺. In such cases, the suffix -ous may be used to identify the lower state and -ic the higher state. Thus:

Fe²⁺ = ferrous ion

Fe³⁺ = ferric ion

Cu⁺ = cuprous ion

Cu²⁺ = cupric ion

This system of denoting the lower and higher valency states of metal cations has, however, given way to the Stock System in which the oxidation number of the metallic element is indicated by Roman numbers in parenthesis placed immediately after the atom concerned.

The FeCl₂ is denoted Iron (II) chloride rather than ferrous chloride and FeCl₃ is denoted Iron (III) chloride rather than ferric chloride.

1.9 Atomic weight

Although it is customary to use the term ‘atomic weight’, the term ‘atomic mass’ is more appropriate. Whilst weight is the force exerted on the body by the influence of gravity, mass is a measure of the quantity of matter in a body independent of gravity.

Historically, oxygen was taken as a standard of mass measurement and the oxygen atom was assigned a value of 16.0000 atomic mass units (amu). On this basis, helium was found to have an atomic weight of 4.003 amu, fluorine 19.000, and sodium 22.997.

However, since the early 1960s, the isotope carbon-12 has been used as a standard and the amu is now defined as being 1/12th of an atom of carbon-12 in which:

1 amu = 1.6605665 x 10^-24 g

In reality the atomic weights based on carbon-12 are in close agreement with those based on natural oxygen.

Table 1.4 compares the atomic weights of the three subatomic particles which show that the atomic weight of the electron is nearly 4000 times less than that of the nucleus and can thus be ignored. In addition, the proton and neutron both have masses close to unity. Since the atomic number is defined as the number of protons in the nucleus, we would expect that the atomic weight would be just over twice that of the atomic number.

This, in fact, is shown in Table 1.5 which lists the atomic weights of various elements together with their atomic numbers.

How about the mass of a molecule?

In fact it’s really quite simple. The molecular weight or mass is merely the sum of the atomic weights of each element.

Thus, the molecular weight of water (H₂O) would be the sum of the atomic weights of the two hydrogen atoms (1 + 1) and one oxygen atom i.e. 1 + 1+ 16 = 18.

Because there are many substances that do not, in fact, comprise molecules (e.g. salts such as sodium chloride) the term formula weight is often used in place of molecular weight and is calculated in exactly the same way. Thus, the formula weight of sodium chloride (NaCl) would be the sum of the atomic weights of sodium (22.99) and chlorine (35.45) i.e. 22.99 + 35.45 = 58.44.

We saw earlier that an ‘atomic mass unit’ (amu) was defined as having a mass of:

1 amu = 1.6605665 x 10^-24 g

Alternatively we could say that there were:

6.022 x 10²³ amu = 1g

The figure of 6.022 x 10²³ is referred to as Avogadro’s number and is termed a mole and is given the symbol mol. A mole is a measure of the amount of a substance and 1 mol of any substance contains 6.022 x 10²³ molecules or atoms.

The formal SI definition of a mole is:

“… the amount of any substance that contains as many particles as there are atoms in 12 g of carbon-12. When the mole is used the particles (atoms, molecules, ions, electrons, etc.) must be stated.”

Thus:

One mole of carbon-12 atoms thus has a mass of 12 g.

One mole of oxygen atoms has a mass of 16 g.

One mole of oxygen molecules has a mass of 32 g.

1.10 Molar concentrations

Many chemical reactions are carried out in solution and the concentration is usually defined in terms of the number of moles of the substance (the solute) contained in 1 ℓ of solution. This is called its molar concentration or molarity.

For example, if 10 g of NaCl is dissolved in distilled water to give 1 ℓ of the solution, what is the solution molarity?

We have already determined that the formula weight of NaCl = 58.44 and thus, by definition:

58.44 g of NaCl = 1 mol

Then:

10 g of NaCl

= 0.171 mol

This is the thus the amount of NaCl dissolved in 1 ℓ of the solution and the molarity is 0.171 mol/ℓ.

1.10.1 Acids and bases

Whilst in a purely covalent bond the electrons are shared symmetrically, in most cases the shared electrons tend to be closer to one atom’s nucleus that the other. Such bonds have a partial ionic character in which one part of the molecule has a net positive charge and the other a net negative charge. These molecules are said to be polar.

Figure 1.31 shows a molecule of water in which the shared electrons are more likely to be found around the oxygen atom. The result is that there is a net negative charge on the oxygen atom. And since the electrons spend less time orbiting the hydrogen atoms there is a net positive charge on each hydrogen atom.

It should be noted that, as distinct from ionic bonding, the net charges are less than 1+ or 1-, this being indicated by the symbol (δ).

In pure water, a small number of molecules ionise — each forming a hydrogen ion (H⁺) and a hydroxide ion (OH^–). This reaction is called self-ionisation. Since the number of hydrogen and hydroxide ions is equal, the water is neutral. However, this balance can be upset by a number of compounds that dissolve or react with the water and produce either hydrogen ions or hydroxide ions.

Substances that react in this manner and produce hydrogen ions (H⁺) in the solution are called acids and those that react and produce hydroxide ions are called bases.

Another name used to describe the hydrogen ion (H⁺) is proton and thus acids may also be described as proton donors and bases as proton acceptors.

In reality, the hydrogen ions do not exist on their own but are attached to the water molecules to become what are called hydroxonium ions (H₃O⁺). This is shown in the reaction of hydrochloric acid (HCl) with water:

HCl + H₂O → H₃O⁺ + Cl^–

Nevertheless, because only the hydrogen ion takes part in the reaction in a solution, the H+ ions may be considered as the active ingredient and, for all intents and purposes, the hydroxonium ion may be considered to be a hydrogen ion.

A base has been described as a substance that gives hydroxide ions (OH^–) when reacting or dissolving in water. Alternatively, it has been described as a proton acceptor. A base is, therefore, the chemical opposite to an acid. As a result, a base may be further described as a substance that will neutralise an acid by accepting its hydrogen ions.

Thus, in the right proportions, the reaction of hydrochloric acid (HCl) with sodium hydroxide (NaOH) gives:

HCl + NaOH → H₂O + NaCl

resulting in a solution containing ordinary table salt.

When a base is dissolved in water the resulting solution contains more hydroxide ions than hydrogen ions and is termed an alkaline. The determination of whether a solution is acidic or alkaline, and their relative strengths, is compared on a pH scale.

1.11 Oxidation-reduction

In our discussions on bonding, we saw how both ionic and covalent bonding involved a shift in electron density from one atom to another.

In the case of sodium chloride (NaCl), for example, an electron was transferred completely from Na to Cl to form the Na⁺ and Cl^– ions (Figure 1.32). This shift in electron density, from one atom to another, is termed an oxidation-reduction reaction — usually shortened to Redox reaction.

Oxidation refers to the loss of electrons during a reaction whilst reduction refers to the gain of electrons. Oxidation and reduction always occur together — with one substance accepting the electrons that another loses. Thus in the reaction of sodium chloride, sodium loses an electron and is therefore oxidised whilst the chlorine atom gains an electron and is said to be reduced.

Although in the sodium chloride reaction, the sodium supplies the electron, and is thus oxidised, it acts as the reducing agent. Likewise the chlorine, which is reduced, acts as the oxidising agent.

A measure of the power of a substance to gain electrons is called the Redox potential. A reducing agent that readily loses electrons will have a negative potential whilst oxidising agents will have a positive potential.

Objectives

When you have completed this chapter you should be able to:

• Describe the basic principle of the electrochemical cell
• Show how a potential difference is formed between two dissimilar electrodes immersed in water
• Demonstrate how a simple voltaic cell is formed
• Explain how a Daniell cell works
• Understand the use of an electrolytic bridge
• Discuss the importance of the electrochemical series and how it is applied

2.1 Introduction

If a metallic rod such as zinc is placed in water, there is a tendency for a redistribution of electric charges to take place.

Firstly, the water molecules attract positive ions from the metal. Countering this, however, is the fact that as the positive ions leave the metal, the rod has a tendency to become negatively charged. This results in the positive ions being held to the surface of the rod — with the ions thus being said to be adsorbed onto the surface of the metal (Figure 2.1).

Some of the adsorbed ions may break away from the surface of the metal and the rod will actually lose mass (Figure 2.2). However, because each ion that is lost causes the rod to become more negatively charged this reaction, called desorption, can only happen to a few ions before an equilibrium is rapidly reached.

The amount of negative charge on such a metal rod placed in water is determined largely by the reactivity of the metal. Zinc atoms ionise easily and thus the negative charge will be relatively high — with a potential difference thus existing across the zinc/solution junction. However, copper atoms, for example, ionise only with some difficulty and thus, as shown in Figure 2.3, the amount of negative charge is considerably less. The result is that the potential difference across the copper/solution junction is also considerably less.

2.2 Potential difference

Assume now that a zinc rod and a copper rod, referred to as electrodes, are both placed in water and a high impedance voltmeter is connected across them (Figure 2.4). As a result of their different electrode potentials, the voltmeter would indicate that a potential difference now exists, between the copper and zinc electrodes, of about 1.1 V.

If, instead of the high impedance voltmeter, a conducting wire is connected across the electrodes there will be an initial brief flow of electrons from the zinc electrode to the copper electrode (Figure 2.5). Since the zinc electrode loses electrons it will become less negatively charged whilst the copper electrode will gain electrons and will become more negatively charged.

The reduced negative charge of the zinc electrode means that the adsorbed positive ions are now less strongly attached to its surface and may become detached — exposing new atoms. Since this encourages more ionisation, the zinc electrode will again tend to become more negative.

At the copper electrode, the inflow of electrons attract adsorbed copper ions even more and further ionisation is thus discouraged, As a result, both electrodes end up at the same potential and no potential difference will now exist across them (Figure 2.6).

2.3 Simple voltaic cell

One means of overcoming this stalemate is to make the solution acidic (e.g. a dilute solution of sulphuric acid) and thus increase the number of free hydrogen ions (H⁺) in the solution. As before, ions will continue to leave the zinc electrode — with ionisation still taking place and the excess electrons being conducted through to the copper electrode.

Now, at the copper electrode, the inflow of electrons attracts the positive hydrogen ions, which are adsorbed onto the electrode’s surface and then combine with the electrons to form hydrogen gas, which bubbles to the surface (Figure 2.7). In this manner, a current will continue to flow through the conductor until either the zinc electrode has been completely corroded away or until all the hydrogen ions in the solution (called the electrolyte) have been released as gas at the copper electrode.

One of the major defects of this simple voltaic cell is that the liberated hydrogen forms a layer on the copper electrode (Figure 2.7). Apart from considerably increasing the internal resistance of the cell, the copper/solution junction now becomes a copper/hydrogen junction that produces an emf that is in opposition to that produced by the current flow. The result is that on connecting the two electrodes via the conductor to produce a current flow, within a very short time, the emf across the electrodes will fall from about 1.1 V to 0.7 V or even less. This effect is termed polarisation.

2.3.1 The Daniell cell

One method of overcoming this problem is by means of the Daniell cell shown in Figure 2.8 — a modified form of the simple voltaic cell. Here the zinc electrode is again placed in a dilute solution of sulphuric acid (or, more traditionally, zinc sulphate). The difference is that in this case the copper electrode is effectively immersed in a saturated solution of copper sulphate which acts as the depolariser.

The cell is redrawn in Figure 2.9 to illustrate the action which shows that, in effect, we now have two half-cells — the zinc electrode in the sulphuric acid forming one half-cell and the copper electrode in the copper sulphate forming the other half-cell. Earlier we had seen how the positive hydrogen ions of the sulphuric acid were used to attract the excess electrons in the copper electrode. In this case, the Cu²⁺ cations of the copper sulphate (Cu²⁺SO₄) are used to attract the excess of electrons in the copper electrode — with the ions thus deposited as metallic copper on the electrode.

2.4 Electrolytic bridge

The two half-cells are connected ‘electrically’ by the porous pot which prevents any rapid mixing of the two electrolytes but at the same time allows ions to diffuse through it.

If the two solutions were not in contact in this way (Figure 2.10) the electrode reactions would rapidly cease. Zinc cations (Zn²⁺) leaving the zinc electrode would rapidly build up a positive charge in the electrolyte and ultimately prevent electrons leaving the zinc electrode. Similarly, the adsorption of copper ions on the copper electrode would leave the solution with an overall negative charge (SO₄^2-) that would prevent negative electrons entering the copper electrodes.

The effect of the porous barrier (Figure 2.11) is thus to allow the negative sulphate anions (SO₄^2-) to pass through the barrier and neutralise the positive zinc cations (Zn²⁺), and vice versa, so that both the solutions would remain neutral.

One of the problems of such a simple porous clay barrier is the fact that different rates of diffusions of the cations and anions can occur across the barrier — giving rise to a liquid junction potential.

This effect can be considerably reduced by using an alternative method of connecting the two half-cells by means of what is termed a salt bridge (Figure 2.12). In this example the salt bridge comprises a glass tube filled with an electrolyte (usually potassium chloride (KCl) or potassium nitrate (KNO₃)) with porous plugs at either end.

These salts produce ions with approximately equal diffusion rates, with either negative anions diffusing from the salt bridge into the copper half-cell or Cu²⁺ cations diffusing into the salt bridge.

2.5 Electrochemical series

The measured potential for a given cell reaction is dependent on the concentration of the ions, the temperature, and the partial pressure of any gases involved in the reaction.

Earlier we saw that it is impossible to measure the potential associated with an individual half-cell and that a potential difference can only be measured when the two half-cells are connected.

In practice the hydrogen electrode is not easy to use as a general standard and is often replaced by the calomel (mercury chloride) electrode, whose potential is defined in terms of the hydrogen electrode.

Despite the fact that the potential of an isolated half-cell cannot be measured, values have been determined by comparing a range of half-cell reactions with that of an arbitrary reference electrode — a hydrogen electrode — which has been assigned an electrode potential or what is termed a standard reduction potential of 0.00 V.

Figure 2.13 shows the standard reduction potentials (otherwise known as the electrochemical series) for a number of half reactions at a fixed temperature of 25°C, a fixed pressure of 1 bar, and a fixed ion concentration of 1 mol/ℓ.

Since the measured cell potential represents the difference between the reduction potential of one half-cell and the reduction potential of the other, this series now enables us to determine the standard cell potential (E⁰_cell ) for any cell from:

E⁰_cell = [Potential of reduced substance] – [Potential of oxidised substance]

Thus, for a zinc/copper cell the standard cell potential would be given by:

E⁰_cell = [+0.34] – [-0.76] = 1.1 V

and for a silver/copper cell:

E⁰_cell = [+0.8] – [+0.34] = 0.46 V.

In the calculations given above, no account has been taken of temperature or the ionic concentration as these were assumed to be fixed at 25°C and 1 mol/ℓ respectively.

The relationship between the cell potential and the concentrations of the reactants are accounted for in the Nernst equation which states:

where:

E = total potential (in millivolts) between two electrodes

E^o_cell = standard cell potential

R = universal gas constant (Joules/mol-Kelvin)

T = absolute temperature (Kelvin)

n = charge of the ion

F = Faraday constant (Coulombs/mol)

a_i = activity of the ion

The significance of this equation is that if all the other factors were held constant, then the emf of the cell will vary according to the ion activity and, therefore, according to its concentration.

Further, if the electrode were made sensitive only to the H⁺ ions, the emf of the cell would indicate the acidity /alkalinity of the medium.

Objectives

When you have completed this chapter you should be able to:

• Show the need for pH measurement
• Describe the properties of water
• Define pH
• Demonstrate the principle of both the measuring and reference electrodes and their relationship to each other
• Understand the Nernst equation and its dependence on temperature
• List the various sources of error in the measurement of pH
• Show how calibration is carried out

3.1 Introduction

The measurement of pH is one of the oldest chemical analysis methods in the world. Using our sense of taste most people can easily determine that orange juice is more acidic than milk and the ‘cola’ based beverages are even more acidic still. Interestingly, what is regarded as ‘good tasting food’ is acidic in nature (Figure 3.1).

In industrial processes the measurement and control of the acidity/alkalinity levels of process media is becoming increasingly more important. Application areas include: neutralisation of effluent in steel, pulp and paper, chemical and pharmaceutical manufacturing; cooling tower control; maximising efficiency in plating and surface treatment; control of municipal drinking water and wastewater purification plants; and quality control in the food and beverage industries.

In its simplest definition, pH could be defined as a measure of the acidity or alkalinity of a solution. However, whilst a 5% solution of boric acid can be used as an eye wash, the use of a 5% solution of sulphuric acid would be disastrous. Thus the knowledge of only the concentration of an acid or base is of little practical use.

One of the most important identifying factors of an acid or base is the hydrogen activity. Both boric acid (B(OH)₃) and sulphuric acid (H₂SO₄) contain hydrogen. However, whilst the hydrogen in sulphuric acid dissociates in the presence of water to become free hydrogen ions, very little of the hydrogen in boric acid is released as free hydrogen ions. Thus, the true measure of acidity (or alkalinity) concerns the measurement of the dissociated or free hydrogen ion concentration of a given solution.

3.2 Properties of water

Earlier we saw that even pure water ionises into an equal number of hydrogen (H⁺) and hydroxide (OH^–) ions. This self-ionisation is very small and at 25°C the molar concentration of each is only 1.10^-7 mol/ℓ.

In all other aqueous solutions, the relative concentrations of each of these ions are unequal such that as one increases the other decreases to form a constant.

This relationship may be expressed as:

[H⁺] x [OH^–] = 1 x 10^-14 = K_w

in which K_w is referred to as the ion-product constant for water.

In essence this shows that the product of hydrogen and hydroxide ions is equal to 1 x 10^-14 (mol/ℓ)². Therefore, if the concentration of one increases by a factor of 10, the other decreases by a factor of 10.

In an aqueous solution where the hydrogen concentration is greater than 10^-7 the solution is acid and if less than 10^-7 the solution is alkaline.

3.3 Definition of pH

pH is derived from the initial letter of the French word ‘potenz’ meaning power and thus expresses the power of hydrogen. pH is actually defined as the negative logarithm of the hydrogen ion concentration and is expressed mathematically as:

pH = -Log [H⁺]

where:

[H⁺] is hydrogen ion concentration in mol/ℓ.

This value ranges from 0 to 14 pH — with values below 7 pH exhibiting acidic properties (an increase in hydrogen ions) and values above 7 pH exhibiting alkaline properties (an increase in hydroxide ions).

Since the pH value is an expression of the ratio of [H⁺] to [OH^–] concentration, then at 7 pH, the ratio of [H⁺] to [OH^–] is equal and the solution is neutral. Because the pH equation is logarithmic, a change of one pH unit represents a 10 fold change in concentration of hydrogen ions. Table 3.1 illustrates the relationship between pH and the H⁺ and OH^– concentrations.

In the former examples, a 5% sulphuric acid solution will have a pH of approximately 0.3 and the 5% boric acid solution will have a pH of about 5 (Figure 3.2).

It might at first appear that the pH scale covers a wide range — from 0 (very acidic) to 14 (very alkaline). In reality, however, the pH scale is designed for diluted acid and alkaline solutions and is of little use in determining concentrated acidic or alkaline species.

From the above table, at pH = 0 (the most acidic on the pH scale), the hydrogen ion concentration, [H⁺], is 1 x 10⁰ or 1 mole.

The term ‘molarity’ describes the concentration of a substance within a solution and is defined as: ‘that concentration which contains that mass of substance which is equivalent to its molecular (atomic) mass in grams dissolved in one litre of water’.

Since the atomic mass of hydrogen is approximately equal to one, a 1 molar solution of acid would contain 1 g of hydrogen ions [H⁺] per litre or that which is equivalent to 1 part in 1000 (1000 ppm) [H⁺] — a relatively low concentration of acid. Table 3.2 illustrates the minimum and maximum ranges of the pH measuring scale.

3.4 Measurement of pH

The measurement of pH in an aqueous solution is most commonly made using a hydrogen sensitive glass electrode that forms one half-cell of an electrochemical cell. The other half-cell comprises the reference electrode that is immersed in the same solution.

The measuring circuit is completed by the meter itself — a high impedance input voltmeter used to measure the emf of the two electrodes (Figure 3.3).

3.5 The measuring electrode

The measuring electrode (Figure 3.4) comprises a glass envelope containing a buffer solution of known ionic strength and pH. A silver wire, coated with silver chloride, is immersed in the buffer solution and forms the measurement lead-off wire.

The sensing element at the tip of the electrode is a hydrogen-sensitive membrane manufactured from special glass containing alkali and alkali-earth and/or rare-earth ions.

pH measurement is based on the chemical reaction that takes place between the sample solution and the membrane surface — generating an electrical potential that varies with the pH value.

The mechanism involved is dependent on hydration of the glass membrane surface which forms a gel layer (hydrated silica) when the electrode is in contact with an aqueous medium. Due to the presence of the encapsulated internal buffer, a gel layer is also formed on the internal surface of the glass membrane.

3.5.1 Hydration

A solution may be defined as a mixture in which the molecules of a substance added to a liquid are evenly dispersed. When the liquid, called the solvent, is water, the action of dissolving the substance is called hydration.

If we dissolve common salt (sodium chloride) in water, hydration occurs as a result of the charged ends of the polar water molecules attracting the ions within the ionic lattice. In this manner, the δ+ charges on the water molecule attract the CL- anions and the δ- charges attract the Na+ cations. As a result the water molecules form cage-like formations around the Na+ and CL- ions — with the negative poles of the water molecules dipole pointing towards the positive Na+ cation (Figure 3.5 (a)) and the positive poles of the water molecules pointing towards the CL- anions (Figure 3.5 (b)).

The forces involved in hydration can be very strong and for many salts, if the water is evaporated, some water molecules remain attached to the ions and actually become part of the crystal. Examples include: plaster of Paris, Epsom salts, borax and washing soda.

As shown in Figure 3.6, singly charged alkali metal cations (e.g. Li⁺ ) in the glass are exchanged for the hydrogen ions in the solution — giving rise to a phase boundary potential at the surface of the hydrated glass layer.

As soon as the activity of hydrogen ions is different in the two phases, diffusion of the hydrogen ion in (acid solution) or out (alkaline solution) of the gel layer occurs. This leads to a build up of charge at the outer phase layer and a new equilibrium state exists which prevents further hydrogen ion transport. The metal ions in the glass membrane are responsible for the migration of the charge to the gel layer on the inner surface of the membrane and it is this potential difference that allows pH measurement to take place.

3.6 The reference electrode

The measuring electrode forms only one half-cell of the pH electrochemical cell and in order to complete the measurement circuit it must be combined with the other half-cell — the reference electrode. In practice, the reference electrode causes more problems than the measuring electrode.

Because the reference electrode provides the reference potential, any deviation in its potential will cause the overall potential to change — thereby causing the overall pH reading to change.

The reference electrode comprises an internal reference element immersed in an electrolyte solution — both housed in a glass or plastic envelope (Figure 3.7). At the tip of the electrode a porous junction, or salt bridge, forms a liquid junction between the electrolyte and the process liquid being measured — thus completing the pH measuring circuit.

3.6.1 Reference element

The most commonly used reference element comprises a silver wire coated with silver chloride. When used in conjunction with a saturated potassium chloride electrolyte, it has a half-cell potential of 199 mV at 25°C.

In some electrodes the silver/silver-chloride wire is immersed direct into the internal electrolyte whilst in others the wire is surrounded by silver chloride crystals within a glass cylinder and separated from the electrolyte by a cotton-wool plug (Figure 3.8).

At one time the calomel electrode was also frequently used. The calomel electrode comprises a silver or platinum wire surrounded by a paste of mercury and mercurous chloride that is again contained within a cylinder sealed with a cotton-wool plug.

The calomel electrode has lost much of its popularity in recent years — due mainly to its use of mercury and the potential danger of contamination to the process media and to the environment.

3.6.2 Electrolyte

In theory any conducting medium could be used as the internal electrolyte. However, the filling solution should fulfil a number of conditions:

• it should have a high ionic strength to minimise resistance
• it should be stable
• it should not react with the process solution
• it should be soluble in the process solution
• the diffusion rates of the cations and anions of the electrolyte salt should be approximately equal.

This last consideration is to prevent the build up of a charge across the liquid junction, which can cause offset errors.

The migration velocity of ions is determined by their charge and size. Table 3.5 shows the ionic mobilities of various ions at 25°C.

Ion	Ionic mobility (cm/s.V)
Li+	4.01 x 10-4
F–	5.74 x 10-4
Na+	5.19 x 10-4
K+	7.62 x 10-4
CL–	7.91 x 10-4
H+	36.25 x 10-4

From this table it may be seen that both K⁺ and CL^– ions have very similar mobilities and thus potassium chloride is the most commonly used electrolyte salt. If, for example, sodium chloride were used, where the mobilities of Na⁺ and CL^– ions are quite different, a positive charge would eventually build up on the internal surface of the liquid junction since the positive Na⁺ cations move slower than the negative CL^– anions (Figure 3.9).

Experience has shown that 3.0 molar KCL solutions fulfil these conditions over a wide temperature range. At higher molar concentrations, the high chloride ion concentration causes the potassium chloride to react, to some extent, with the silver chloride of the internal reference element to form soluble complexes. This can be overcome by saturating the potassium chloride electrolyte with silver chloride — thus providing a filling solution that is in equilibrium with the internal electrode. Such an electrolyte is not recommended for use with process solutions containing sulphide ions and sulphide compounds, cyanide ions and other halide ions, because of the precipitation of highly insoluble silver compounds.

Potassium chloride itself cannot be uses as an electrolyte with all process solutions since reactions can take place. Examples include: the presence of mercury (II), silver, copper (I), and lead (II) ions that react with chloride ions to form insoluble compounds.

3.6.3 Reference junction

The reference junction, also referred to as a ‘salt bridge’, liquid junction or ‘frit’, forms the interface between the process solution and the internal electrolyte and completes the electrical path from the measuring electrode to the reference electrode.

Apart from allowing small amounts of electrolyte to flow through it, the reference junction should also be chemically inert so as not to interfere with the ion exchange process. It should also maintain a low consistent electrical resistance value.

Usually, the junction comprises a ceramic frit formed in the shape of a cylindrical plug about 2-3 mm in diameter and between 3-12 mm in height. However, use is also made of a variety of porous materials including: wood, Teflon, kynar, asbestos and quartz fibres. An alternative, the ground glass sleeve junction, is mainly used for laboratory purposes and comprises two ground glass areas, mated to each other, that allow the electrolyte to permeate between them. Annular junctions, which surround the measuring electrode, are also used. Some electrodes incorporate a replaceable junction using an O-ring as the seal.

In order to prevent the process solution from flowing into the reference and contaminating the KCL solution, and/or attacking the Ag/AgCL reference element, the junction is designed to allow small amounts of the electrolyte solution to leach through it into the process.

The electrolyte housing, therefore, not only forms a reservoir for the electrolyte that is leached away, but also provides sufficient head to maintain a positive pressure. Where high process media pressures exist or high process media temperatures could produce reversed flow across the liquid junction, this pressure could be insufficient. Consequently, the filling hole is often connected to a further electrolyte reservoir or pressure source (Figure 3.10).

Many of these problems are overcome in modern electrodes that make use of a gelling agent that is added to the electrolyte solution. Apart from slowing down the effects of contamination through the porous junction the gel-fill may be sealed with a positive pressure since it is less susceptible to being forced out of the junction during periods of heat or pressure cycling.

3.6.4 Reference construction

In practice, the measuring and reference electrodes are usually combined into a single unit (Figure 3.11) in which the reference electrode is arranged concentrically around the measuring electrode.

The combination electrode is much easier to handle than separate electrodes and is, therefore, used almost exclusively in industrial applications. Only when the two electrodes are expected to have very different life expectancies is the use of separate electrodes recommended.

Apart from preventing contamination of the internal environment of the reference electrode, it is also important to prevent coating forming over the junction.

The liquid junction can be blocked if the stream contains any material that reacts with the filling solution to form a precipitate. Particularly troublesome are silver, lead and mercury, which form insoluble chlorides.

This may be overcome using the double-junction reference electrode which is essentially a complete electrode with its own liquid junction fitted within an electrode outer body having a second liquid junction in contact with the sample. The main advantage of this electrode is that the reference solution in the outer body, usually potassium chloride, can be chosen to be compatible with the ‘inner’ electrode solution and the sample. Clogging of the outer junction is extremely unlikely, since neither potassium ions nor chloride ions form insoluble compounds with the majority of materials found in process streams. Blocking of the inner junction is not possible.

3.7 Nernst equation

The Nernst equation relates the cell potential with the hydrogen ion concentration:

where:

E = potential generated (in millivolts)

E^o = standard cell potential

R = universal gas constant

T = absolute temperature (K)

n = valency (charge) of the ion (n = 1 for H⁺ ion)

F = Faraday’s constant

[H⁺] = hydrogen ion concentration

The term is called the ‘Nernst’ or ‘Slope’ factor and is given the symbol N and the equation becomes:

E=E° + N log[H+]

Thus at 25°C if the pH changes from 7 to 8 the voltage output will change from 0 to 59.16 mV. Similarly a change in pH from 8 to 9 will result in a voltage change from 59.16 to 118. 32 mV.

Table 3.6 shows the expected mV output for a number of pH values at 25°C, whilst Figure 3.12 illustrates this.

3.7.1 Temperature effect

It should also be very apparent that the ‘Nernst factor’ also determines the number of millivolts for each pH unit at different temperatures. This is shown in Table 3.7.

Temperature (°C)	mV/pH decade
5	55.19
10	56.18
15	57.17
20	58.17
25	59.16
30	60.15
35	61.14
40	62.14
45	63.13
50	66.12
55	65.11
60	66.10
65	67.10
70	68.09
75	69.08

In effect, the effect of temperature is to change the slope of the response as shown in Figure 3.13. Since the change in output vs. temperature is linear, it can be expressed as 0.00335 pH error/pH unit/°C.

Thus, if an uncompensated pH system were standardised in a pH 7 buffer at 25°C, then a sample that measured pH 3 at 15°C, would have an error of:

0.00335 x 10°C x 4 units = 0.134 pH unit

And at 75°C (probably close to a typical worst case), a 3 pH sample would read 3.67 pH using an uncompensated pH system.

Clearly, then, the measurement of pH must also involve the measurement of temperature in order to compensate for the errors due to the Nernst slope.

3.8 Antimony electrode

When measuring effluents containing hydrofluoric acid (HF), which attacks glass, use is often made of the antimony electrode. Used with a conventional reference electrode, the antimony electrode consists of a plastic body with an antimony ring fitted into the end (Figure 3.14).

When first immersed in a solution containing oxygen, the surface of the antimony oxidises to form the pH responsive layer. The electrode also usually includes a built-in scraper or grinding mechanism that is used to restore the electrodes performance should it become coated during use.

Because the antimony electrode is extremely rugged, it is often used for determining the pH of soils. Further, because antimony is extremely hard, it may be used in high flowing abrasive liquids (slurries ) that would normally grind electrodes away.

However, antimony ring electrodes have many disadvantages:

• since it is a metal electrode it will respond to Redox potentials;
• any metal more noble than antimony i.e. having a standard half-cell potential greater than antimony, will deposit on the electrode;
• at pH = 7 the potential of the antimony measuring chain is approximately -400 mV and compensation within the circuitry of the transmitter must be made to accommodate this potential shift;
• the slope of the antimony measuring chain differs to that of the glass electrode 52 to 57 mV/decade;
• the response is close to Nernstian only over a narrow range and thus temperature compensation is only effective over the pH range 2 to 7;
• antimony and antimony oxide are poisonous.
• the best reproducibility normally attained is about 0.1 pH
• response time is poor (from 3 to 30 minutes)

Because there is so much more uncertainty to every aspect of behavior of the antimony electrode compared with the glass electrode, it should be considered only as a last resort.

Even when working with hydrofluoric acid it should be noted that HF solutions above 4 pH do not have such detrimental effect on glass electrodes. Under these circumstances it may be better to use a glass electrode offering superior performance despite a relatively short life.

3.9 Sources of errors

The measurement of pH is subject to a wide number of potential errors and in order to ensure accuracy of measurement, each must be taken into account.

3.9.1 Solution temperature coefficient

While temperature compensation is applied for changes in the electrode pair output, both in terms of slope and zero offset, it should be remembered that most solutions change their pH value with temperature variation (Figure 3.15).

For some applications, it is possible to refer the value to a reference temperature, e.g. 25°C, so removing the effect on the measurement of temperature change. Compensation has also been incorporated into the design of several industrial pH meters.

3.9.4 Alkaline error

Alkaline error (also referred to as sodium ion error) is the result of sodium ions, partly or completely, replacing the hydrogen ions — thus decreasing the hydrogen ion activity and artificially suppressing the true pH value.

Sodium ion interference occurs when the hydrogen ion concentration is very low and the sodium ion concentration is very high. The effect thus increases with increasing pH value (pH > 9) as well as increasing temperature (Figure 3.16).

Since there is no glass formulation currently available with zero sodium ion error, it is important that the error be consistent and repeatable. Such characteristics are available when controlled molecular etching is applied to the glass membrane. This provides a consistent amount of lithium ions that are available for exchange with the hydrogen ions to produce a similar milli-volt potential for a similar condition.

3.9.6 Glass electrode impedance

Whilst glass is usually considered to be a near perfect insulator, in reality the glass electrode is an ionic conductor whose resistance changes according to the solution temperature — decreasing as the temperature rises.

Because the pH responsive membrane is in series with the electrical circuit it should have as low a resistance as possible whilst combining good pH response, durability and mechanical strength. These requirements cannot be achieved with a single type of glass membrane for all applications. As a result, three different types of glass electrode are produced (Figure 3.17).

The ‘standard’ membrane glass, used for the majority of purposes; has a good response over 0 to 14 pH and can be used between 0 and 100°C. The ‘low resistance’ membrane glass has a much faster pH response by a factor of 10 (because of its lower resistance) but can only be used over a pH range of 0 to 10 and up to temperatures of 60°C. And the ‘high temperature’ membrane glass is used at temperatures between 50 and 130°C.

Consequently, the resistance of the membrane glass can lie between 1 and 1000 MΩ depending on the process temperature and the type of electrode selected. In order to develop most of the electrode potential across the input impedance of the pH meter, and to minimise the effect of impedance variations due to changes in temperature, the pH meter input impedance needs to be at least 1000 times higher than the maximum electrode impedance of 1000 MΩ, i.e. 10¹² Ω (Figure 3.18).

Fortunately, modern high input impedance amplifiers are now available in the range 10¹³-10¹⁴ Ω. Thus, with an amplifier having an input impedance of 10¹³ Ω, for example, the error due to impedance matching is less than 0.01% and a change in electrode impedance from 1 to 1000 MΩ will effect the voltage by only 0.001%.

The main implications of the very large electrode source impedance lies in maintaining the high insulation in all electrode terminations. This insulation has to be maintained, not only between the glass and reference terminations, but also between glass and earth, since the solution is often at earth potential.

The most common problem encountered, especially in damp conditions, is low insulation of the electrode terminations. In addition, replacement components must be carefully chosen in order to obtain a high measurement integrity, since construction materials used in many plugs and sockets and termination blocks are totally inadequate for pH use due to their low insulation.

Although the insulation of the electrode terminations cannot be measured directly, due to the lack of commercially available insulation testers, the presence of low insulation can be determined using a pH electrode simulator. This device, a vital piece of test equipment for any fault finding investigations on pH systems, is basically a mV source with the added feature of being able to switch in a 1000 MΩ resistor into the glass electrode output lead to simulate the impedance of the electrode.

Calibration

In the pH vs mV curve of Figure 3.19, we can see that there are two deviations that may occur from the ideal relationship at a given temperature: zero shift and slope deviation.

Zero shift occurs when the reference electrode no longer performs correctly or if the pH electrode itself is fouled. The effect is to move the curve along the x axis i.e. the potential is no longer 0 mV at pH = 7 but at, for example, pH =6 or 8. Most analysers allow an adjustment of ± 2 pH units.

The effect of slope deviation is to alter the slope from 59.16 mV/decade. Because it is not possible to always manufacture electrodes with a 100% response a slope adjustment of 95 to 105% is usually allowed in most instruments. Other factors affecting the slope are due to faulty manufacture, ageing of the electrode or fouling of the membrane.

Since both these errors must be catered for, two-point calibration is usually performed using a buffer solution. A buffer solution is a pH solution that maintains a nearly constant pH if diluted or concentrated and despite the addition (possibly due to contamination) of small amounts of acid or alkali. Standard buffer solutions are available from instrument suppliers and normally span the following ranges:

‘pH 1’	DIN 19267	0.1 M hydrochloric acid.
‘p H4’	BS1647	0.05 M potassium hydrogen phthalate.
‘pH 7’	Mixture of disodium hydrogen phosphate and monopotassium dihydrogen phosphate.
‘pH 9’	BS 1647	0.05 M disodium tetraborate (borax).
‘pH 10’	BS1647	0.025 M sodium hydrogen carbonate + 0.025 M sodium carbonate.

Normally the zero shift is set first using a buffer with pH = 7 (Figure 3.20(a)). Next, the slope is set using a buffer with a 2 to 3 pH unit difference from the first (Figure 3.20(b)). Since the electrode response may not be perfectly linear across the entire 0-14 pH range, a pH buffer with a value greater than 7 (e.g. pH = 9 or 10) is used when measurements above 7 are carried out and vice versa.

In some cases, ‘bracketing’ of the measured value may be necessary in which two buffers, one below and one above the average measured value, are used. This ensures greater accuracy over the measurement interval.

3.10 Applications

3.10.1 High-purity water and pH (Courtesy: MILLIPORE APPLICATIONS Technical Library)

pH influences the rate of many chemical and enzymatic reactions. Therefore, buffered solutions are commonly utilized in the laboratory. These solutions are usually prepared with high-purity water in order to minimize the risk of contamination with water impurities. Discussed here are the difficulties and alternatives to measuring the pH of high-purity water and the impact of buffer dilution on pH.

pH of high-purity water

The pH of water is defined as the negative logarithm of the hydrogen ion activity, which is often assimilated to the hydrogen ion concentration (Equation [1], Table 3.8). While the pH of pure water is approximated to be 7.0 at 25°C, the theoretical value is 6.998. Indeed, the dissociation of water and the ionic product for water (Equations. [2] And [3], Table 3.8) indicate that 10^–6.998 mol/L of H₃O⁺ are present in pure water. However, one cannot measure the pH of high-purity water simply by dipping the electrodes of a pH meter into a beaker full of water. This measurement is challenging due to the quasi-absence of ions that would enable electron transport between the measuring and reference sides of the pH electrode. It often results in erratic and meaningless pH readings.

Adding the slightest amount of acid or base to pure water will change its pH significantly. Pure water readily absorbs carbon dioxide (CO₂) when exposed to the atmosphere, which forms carbonic acid (Eq. [4], Table 3.8). Carbonic acid dissociates into bicarbonate that is in equilibrium with carbonate. Dissolution of CO₂ in water ultimately leads to a pH of approximately 5.8.

It is this pH value that is eventually obtained if one measures the pH of ultrapure water without working under controlled conditions.

The issues linked to the pH measurement of ultrapure water are addressed in various references. The United States Pharmacopeia (USP) recommends the addition of potassium chloride (KCl) to alleviate the problem. A more precise and complex experimental setup and method for the pH measurement of pure water are described in an American Society for Testing and Materials (ASTM) Standard Test Method. It includes the use of a flowthrough sample chamber and the addition of KCl via a reference electrode with positive electrolyte leakage. Other standards, such as the International Organization for Standardization (ISO) 3696 standard, simply do not require the pH measurement of pure or ultrapure water. They rely on other physicochemical parameters to characterize water quality, such as conductivity.

Conductivity of high-purity water

A conductivity measurement may be used as an alternative to the pH measurement of high-purity water. Conductivity measures the flow of electrons through a fluid, which is proportional to the concentration of ions, their charge, and mobility (Equation. [5], Table 3.8). The relationship between resistivity (the reciprocal of conductivity) and the pH of water is illustrated in Figure 3.21. When only pure water is present, a conductivity value of 0.055 S.cm–1 at 25°C is measured. This results from the water dissociation into hydroxide and hydroxonium ions. Therefore, at 25°C, a conductivity of 0.055 S.cm^–1, or a resistivity of 18.18 MΩ.cm, implies that the water is ultrapure and that the pH is inherently 6.998.

Diluting buffer solutions

The pH of a buffer solution composed of a weak acid and its conjugate base is related to the dissociation constant (pK_a) of the acid/base pair, as shown in Equation.[6], Table 3.8. By definition, buffers are resistant to large changes in pH. However, in the laboratory setting it becomes evident that diluting a buffer with pure water impacts the pH of the solution. Several parameters might explain these pH variations.

1. Temperature. Activity coefficients and equilibrium constants vary with temperature, which causes the pK_a and pH of a buffer to change. Temperature coefficients vary significantly with the nature of the buffer (–0.028 pH/°C for Tris and –0.0028 pH/°C for phosphate). For this reason, temperature is usually reported along with pH values. The pH of water itself varies with temperature: 7.27 at 10 °C, 7.00 at 25 °C, and 6.63 at 50 °C.
2. Buffer capacity. Buffer capacity is the ability of a buffer to resist changes in pH. It depends on the total concentration of the acid/base pair and the concentration ratio. Buffer capacity is high when solutions are concentrated, the ratio of acid to base is close to 1, and the pH is ±1 unit from the pK_a.
3. Ionic strength. Ionic strength describes the amount of electric charges in a solution. It is a function of the concentration and the valence of the ions present. High ionic concentration tends to limit the mobility of the hydrogen ion, which decreases its activity and influences the pH measurement.
4. Dilution. Related to a decrease in ionic strength, dilution directly impacts the pH of a solution. For example, the dilution of a 0.1 M buffer system (HA/HA^–) with an equal volume of water results in a pH change of 0.024 units. The pH variation of HA^–/A2^– buffer systems, such as phosphate, is three times that value.

Conclusion

High-purity water is often used in the laboratory to avoid possible artifacts caused by water impurities. Measuring the pH of this water is challenging and necessitates specialized equipment. However, using water with high resistivity (18.2 MΩ.cm at 25 °C) ensures that the pH is close to 7.0. Diluting buffer stock solutions with high-purity water leads to pH changes. These variations in pH are unavoidable and can be explained by changes in ionic strength and buffer capacity.

3.10.2 pH and redox sensors

Some of the sensors from the industry and their features are discussed here.

(From METTLER TOLEDO)

Monitoring of the parameters pH and redox (ORP) is most important in many different industry processes to reduce cost and optimize the yield. The requirements of many industry processes are different. In chemical processes the material has to withstand acids or caustic strong solutions under for example high temperatures. The electrodes must also be certified for use in hazardous areas.

In pharmaceutical and food processes measuring systems have to meet hygienic design, and must withstand many sterility cycles and provide traceability for validation purposes. pH measurement in high purity water requires special precautions. The high purity pH assembly uses a shielded flow chamber and a permanent outflow of reference electrode for stable measurement and a 3-way valve for in-line calibration.

3.10.3 SevenMulti™ Version 2.0

This model of the pH measurement unit from Mettler Toledo supports the following features.

• Increased sample throughput through automated dual-channel GLP measurement and calibrations using the Rondolino sample changer. A barcode reader can be connected to read in the sample ID.
• PIN number based access control avoids accidental changes to the measurement settings.
• The LabX software provides direct connectivity with Microsoft Excel for data transfer.

3.10.4 InPro4501VP

InPro4501VP is a rugged pH-electrode with PVDF shaft, built-in temperature sensor and Solution Ground with diagnostics for chemical applications and high flow-through rates. Also available with fixed cable.

Features and benefits

Low maintenance: A solid XEROLYT®PLUS polymer electrolyte (no refilling and pressurizing) in direct contact with the process media via an open junction. No clogging in contaminated media due to the open junction.

Reliable measurements in heavily contaminated liquid media: It is suitable in industrial processes or in wastewater applications with highly contaminated solutions, emulsions and suspensions, sulfide-bearing media, suitable also for high flow-through rates. The solution ground assures accurate measurements with a built-in sensor diagnostics tool.

Rugged PVDF body design: Dual 1″ NPT threads allow the sensor to be screwed directly into immersion tubes, pipes, process vessels, etc., due to the polymer body design, protecting the electrode from mechanical damage, acting as a housing at the same time.

Flat membrane for abrasive media: The electrode has a flat glass membrane with protective stand-offs, for media with high content of fibers, particles, etc.

Specifications

Short description	low-maintenance pH-electrode with PVDF body
pH-range	1…14
Temperature range	0…100 °C (32…212 °F)
Pressure resistance (bar)	0…7 bar up to 65 °C, 3.5 bar at 100 °C
Pressure resistance (psi)	0…100 psi up to 150 °F; 50 psi at 212 °F
Diaphragm	open annular junction
Reference system	Argenthal system
Solution Ground	yes, with diagnostics
Sterilizable	no
Autoclavable	no
Shaft material	Polymer PVDF, diameter: 28.5 mm or 1.12″
Connector	VarioPin (VP) IP68 resp. fixed cable
ISM	no

3.11 Troubleshooting

Objectives

When you have completed this chapter you should be able to:

• Describe the basic principle of conduction through ionisation and list the factors that effect ionic mobility
• Specify the requirements for cell design and construction
• Determine the need for temperature compensation
• Understand the special problems associated with the measurement of high purity water
• Demonstrate the requirements for good installation practices
• Examine maintenance and troubleshooting procedures
• Discuss various applications

4.1 Introduction

Conductivity is a measure of the ability of a solution to carry an electric current. Conductivity measurement is limited to aqueous solutions containing acids, bases and salts, in which the molecules dissociate to an ionic state. Unlike pH, which only measures the hydrogen ions, conductivity measurement provides information about the total ionic content of a solution.

The main application of conductivity measurement lies in determining the purity of water or the concentration of chemical solutions. In water purifying systems it helps to monitor the level of water treatment needed to avoid costly corrosion damage in boilers and cooling towers. And in the food and pharmaceutical industries it guarantees the purity of makeup water and helps to monitor equipment cleaning processes.

The resistance of the cylindrical solid conductor shown in Figure 4.1, is determined by the length L, the cross sectional area A, and the specific resistance ρ:

The specific resistance ρ is a constant for the material and is usually expressed in ohms.m or ohms.cm.

The reciprocal of resistance R is called conductance G and is defined as:

where κ is the conductivity expressed in Siemens per metre (S/m).

Although the unit of conductivity is S/m it is more common to use Siemens per centimetre (S/cm) where:

1 S/cm = 1 000 mS/cm = 1 000 000 µS/cm

In contrast to electron flow in metallic conductors, the current flow in electrolytes is due to the transport of ions where, generally, the conductivity increases as the concentration increases. As a result, conductivity is a more straight forward parameter to use in the measurement of ion concentration than the reciprocal resistivity.

Figure 4.2 compares resistivity and conductivity over a wide range of chemical processes and applications and shows how with ultrapure water, the resistivity (Ω.cm) is often used in which:

10 MΩ.cm = 0. 1 µS/cm

and

1 MΩ.cm = 1 µS/cm

4.2 Conduction through ionisation

If two electrodes, immersed in a solution of sodium chloride, are connected to a current source (Figure 4.3), a definite sequence of events will occur. A solution may be defined as a mixture in which the molecules of a substance (called the solute) when added to a liquid (the solvent) are evenly dispersed.

1. Positive sodium ions will begin to migrate toward the negative electrode, while negative chloride ions will travel toward the positive side.
2. The sodium ions will accept electrons at the surface of the negative electrode to form elemental sodium. This free sodium will immediately react with water to form NaOH and H₂. At the same time, at the surface of the positive electrode, chloride ions will donate electrons to form elemental chlorine. The net result is that as electron e₁^– enters the negative electrode, electron e₂^– exits the positive — thus creating a unidirectional current in the external circuit.
3. A layer of negative ions will form near the surface of the positive electrode and positive ions will group at the negative electrode. This presents a “double charge layer” effect that acts as a capacitance.

The current that flows in this circuit is initially a function of the applied voltage and, assuming the resistance of the external wires to be negligible, the resistance of the solution. Solution resistance is determined by the nature of the solvent (a fairly constant factor when considering only water), the number of carriers (ions) available, and the mobility of these carriers.

Generally speaking, it is thus possible to measure the current flowing through the circuit and obtain the solution resistance. With this information, the chemical concentration can be determined by comparing the current values with those previously obtained at known concentrations.

Unfortunately, because of a variety of physical and chemical considerations that create interferences, this simple illustration provides reliable information only in very special circumstances. To be practical, a workable system must take into account all of these interferences.

First consider the formation of elemental materials at each electrode (electrolysis). In the example given, gaseous hydrogen and chlorine are present and begin a secondary reaction with the electrode material. This creates a resistance or possibly a source of voltage that is not related to the measurement of interest.

Even if the freed elements were non-reactive, their physical presence would still cause additional resistance and, therefore, interference (gas obviously will not conduct at normal voltages). Also, a sustained voltage at a given electrode can deplete the immediate vicinity of carriers at high concentration, giving a falsely high resistance reading. Dropping the applied voltage level only delays this action slightly — rendering the system more sensitive to interference from external electrical noise and increasing the adverse effects of the double layer capacitance.

One widely used solution to these problems is to use an alternating current (AC) source so that current direction is reversed before electrolysis or carrier depletion becomes pronounced.

4.3 Ionic mobility

Whilst the use of an ac current source will do much to reduce interference caused by electrolysis, other important factors must also be taken into account.

Individual ions are far larger than electrons and, in moving counter-current to their oppositely charged companions, produce a drag effect. This effect is heightened by the fact that each ion is surrounded by ions of opposite charge and the resultant ‘atmosphere’ must be maintained despite the continual and opposite motions of all components involved. Additionally, the ion will be hydrated, which means that it must carry from 1 to 5 water molecules with it, which increases its mass considerably.

These effects increase in severity as chemical concentrations rise and the ionic atmosphere becomes more crowded. The result is a non-linear but repeatable relationship between conductivity and concentration. Direct compensation for this effect using external circuitry is, in most cases, impractical and conductivity meters normally incorporate built-in calibration curves or look-up tables.

For many chemicals the conductivity actually reaches a peak at some intermediate concentration and decreases beyond this point — with one resistance value thus representing two concentrations (Figure 4.4). Obviously, care must be taken in using the measurement in these cases and a separate check is needed to reveal whether the high or low concentration is being measured.

4.4 Practical conductivity measurement

Figure 4.5 shows how the measurement of conductivity is applied — with the geometry given by the length L of the liquid column between the plates and the cross sectional area A of the liquid column.

An electrolyte’s conductivity depends on:

• the nature of the solvent;
• the nature of the ions; and
• the temperature.

To compare conductivity of different electrolytes the term equivalent conductivity λ is introduced, which leads to the equation:

where:

Z_j = charge of the ion j

C_j = concentration of ion j

λ_j = equivalent conductivity of ion j

This indicates that conductivity increases as the ion concentration increases. Indeed, over a certain range conductivity versus concentration is a linear function. However, many strong electrolytes start to show a decrease in conductivity, after they have reached a given concentration.

4.5 Cell design and construction

In practice, conductivity is measured with a conductivity cell whose mechanical properties (L/A) remain constant over the entire temperature range.

In a traditional cell having 1 cm squares of platinum 1 cm apart (Figure 4.6), the cell constant is 1 cm^-1 and its conductance in µS is numerically equal to its conductivity in µS/cm. Such a cell provides a working range of about 10 Ω to 1 MΩ — corresponding to a conductivity range of 0.1 S/cm to 1 µS/cm respectively. However, as illustrated in Figure 4.2, the full range of conductivity measurement extends from 0.05 µS/cm to 1 S/cm.

This problem is overcome by changing the cell geometry (the path length and the cross-sectional area) whilst still maintaining the measurement based on 1 cm. This gives rise to what is called the Cell Constant (K) which is determined by the formula:

K = L/A

As shown in Table 4.1, the total range can be accommodated by offering a range of cell constants. For example, if the spacing is reduced by a factor of 10 (i.e. L = 0.1 cm) then the cell constant becomes 0.1 cm^-1 and the equivalent resistance for a conductivity of 0.1 µS/cm becomes 1 MΩ. And by reducing the spacing by a factor of 50 (i.e. L = 0.05 cm) then the cell constant becomes 0.05 cm^-1 and the equivalent resistance for a conductivity of 0.05 µS/cm becomes 1 MΩ.

Conversely, for high conductivity solutions, increasing the spacing by a factor of 10 (i.e. L = 10 cm) changes the cell constant to 10 and, therefore, the equivalent resistance for a conductivity of 1 µS/cm becomes 10 Ω.

Generally cell constants of 0.1 or 0.01 cm^-1 are used for low conductivity solutions, and sensors with cell constants of up to 50 cm^-1 are available for high conductivity solutions.

It should be noted that in the definition of the basic cell of Figure 4.6 having two 1cm² electrodes separated by a distance of 1 cm, current flow is assumed to occur entirely within the 1cm³ of solution between the electrodes. However, this is only true if the four sides of the cube between the electrodes are enclosed with a non-conducting material. In practice, when the two plates are close together most of the current will, indeed, flow in the desired 1 cm³ volume, but not all of it.

This effect is exacerbated at higher cell constants and compensation must be applied in the design of the cell.

Figure 4.7 shows a typical concentric 2-electrode sensor whilst Figure 4.8 illustrates the construction of a typical in-line electrode.

4.6 Temperature compensation

Because the conductivity of most solutions increases with increasing temperature (Figure 4.9) the sensors normally include a means of measuring temperature.

The conductivity per °C is expressed by the temperature coefficient α. When this coefficient is known for an electrolyte solution the conductivity value can be calculated at any temperature:

κ₁ = κ₀ [1 + α₀ (T₁ – T₀)]

where:

κ₁ = conductivity at T₁

κ₀ = conductivity at T₀

α₀ = temperature coefficient at T₀

T₁ = measured temperature

T₀ = reference temperature

Weak, aqueous solutions (conductivity between 10 and 15000 µS/cm) generally have temperature coefficients of conductance of very nearly 2%/°C. At moderate strengths acids, generally, exhibit temperature coefficients of between 1 and 1.6%/°C, alkalis between 1.8 to 2.2%/°C, and salts around 2%/°C.

4.7 Conductivity measurement of high purity water

High purity water with a conductivity of less than 10 µS/cm may be considered to be made up of two components: the ‘solute’ component and the ‘solvent’ component. The ‘solute’ component is the contribution to the conductivity due to dissolved solids whilst the ‘solvent’ component is the portion that contributes to the conductivity that has no dissolved solids. Water with no dissolved solids has a conductivity of 0.055 µS/cm at 25°C.

These ‘solute’ and ‘solvent’ components each have distinct temperature compensation coefficients whose interaction depends on both the conductivity of the water and the temperature. Figure 4.10 shows that this dependency is non-linear — with the temperature coefficient varying from 7.4%/°C at 0°C to 2.3 %/°C at 100°C.

In order to achieve full automatic temperature compensation, these two components must be compensated for separately, according to the following expression:

where:

G₂₅ = conductivity at temperature 25°C

G_t = conductivity at temperature t°C

G_upw = ultra-pure water conductivity at temperature t°C

α = impurity temperature coefficient

0.055 = conductivity in µS/cm of ultra-pure water at 25°C.

In order to cater for these non-linear dependencies, temperature compensation is normally built into the conductivity transmitter (see Figure 4.7) enabling users to select from a number of different modes: e.g. ultra pure water with traces of NaCl (ASTM #D1 125-9 1); ammonia; NaOH; and natural waters to DIN 38 404.

4.8 4-electrode sensor

Two-electrode sensors are mostly used in pure and clean water applications with a conductivity of < 1 mS/cm. In applications with higher conductivity, however, polarisation and fouling again become a problem. One solution is to make use of the 4-electrode sensor which has been developed to overcome the problems associated with traditional cells.

In the four-electrode sensor shown in Figure 4.11, an alternating current source is connected across the two outer electrodes which produces a current flow through the measuring cell. Located within the electrical field of these two ‘current’ electrodes are two inner ‘voltage’ electrodes that are used to measure the resultant voltage. As shown, the current source is then regulated electronically to maintain a constant voltage at the inner ‘voltage’ pair of electrodes.

As the conductivity changes, the change of solution resistance attempts to change the voltage between the inner ‘voltage’ electrodes. However, the automatic regulation circuitry adjusts the current between the ‘current electrodes’ to reinstate the predetermined constant voltage between the ‘voltage electrodes’. In this manner, the current through the ‘current electrodes’ is proportional to the conductivity.

Apart from polarisation compensation, the four-electrode system also compensates for the effects of fouling (up to a certain level) by automatically compensating for the increase in resistance on the electrode surface.

If the ‘voltage electrodes’ become fouled, the measurement is virtually unaffected because the voltage across them is measured by a very high input impedance amplifier which draws practically no current. The voltage drop across the fouling is, therefore, virtually zero. Further, if the ‘current electrodes’ become fouled, the voltage across them is increased automatically to maintain a constant current — within predetermined limits.

4.9 Installation

Although electrolytic conductivity measuring cells are rugged devices, they are, nonetheless, precise, accurately manufactured instruments that form an integral part of a sensitive measuring system. As a result they should be handled with care and installed in positions where they are adequately protected from mechanical shock and the extremes of temperature.

Basic installation requirements should include:

• avoidance of dead ends, pipe stubs or locations where there is poor flow through the sensor (Figure 4.12);
• flow at point of measurement should not be great enough to cause distortion to the electrodes;
• amount sensor so flow enters open end of sensor;
• sensors can be mounted at any angle as long as there is good flow through the sensor;
• sensor must be fully immersed;
• sensor should be mounted to allow for easy removal;
• avoid mounting where there could be air pockets; and
• avoid mounting where sediment could settle in them.

4.10 Maintenance

4.11 Applications

4.11.1 Reverse osmosis

When solutions of different concentrations are separated by a semi-permeable membrane, the natural flow of the more dilute to the more concentrated solutions is called osmosis. In water purification by reverse osmosis (RO), sufficient pressure is applied to the concentrated side to force the water to flow through the membrane in the opposite direction. Since solid particles cannot easily pass through the membrane, the water is purified, with 85 to 98% of dissolved salts being removed in one pass.

The efficiency of this process is usually monitored by measuring the conductivity of both sides of the membrane in order to determine the percent rejection. A change in the percent rejection would indicate a problem in the efficiency of the membrane system. Figure 4.13 shows a typical RO operation system using a conductivity transmitter with dual sensor input.

4.11.2 Ion exchanger process

Ion exchange systems, also called deionisers or demineralisers, purify water close to absolute purity by removing ions from raw water. Two basic processes are involved: a cation exchange resin removes the positive ions such as calcium, magnesium or sodium, and the anion exchange resin removes negatively charged ions such as sulphate or chloride.

A typical industrial system (Figure 4.14) uses a cation bed, an anion bed and a mixed bed in series. The resulting ultrapure water has a conductivity down to 0.055 µS/cm or a resistivity of 18.18 MΩ.

Over time, the resins used in purification systems become saturated with ions and the water will no longer be purified. Once the resin reaches its limit the resin must be regenerated with a strong regenerating solution. This regeneration process includes backwash, regeneration, rinsing and service.

The measurement of conductivity is used to indicate when regeneration is required and thus reduce the amount of chemicals used and extend the useful life of the ion exchanger resins.

4.11.3 Porosity test using high voltage

To protect objects made of steel or metal, they are coated with corrosion resistant materials such as rubber, synthetics or enamel. The protective coatings must be tight, that is, free of pores cracks or embedded foreign objects, to keep aggressive materials from the carrier material that is in danger of corrosion. Fine pores or cracks cannot be entirely avoided with any coating process.

Measurement method in poroscope * HV20 from FISCHER

The test method is based on the fact that all electrically insulating coating materials have a much higher disruptive strength than air. Pore detection occurs at the defective spots through a spark-over (short circuit) between the test electrode and the conducting base (Figure 4.15). A defective spot may be a thin air channel (pore, crack) or a coating that is too thin over the conducting base underneath.

Features of the HV20 include:

• Low-energy, and therefore safe, high voltage.
• High voltage generation in the test head.
• Two test head versions with test voltage ranges: 0.3 to 3 kV and 2.5 to 25 kV can be connected.
• Extensive electrode selection
• Continuously adjustable test voltage.
• Display of the test voltage that is present directly at the electrode.
• Electronic test voltage monitoring.
• Optical indication at the test head and the test instrument when a pore is detected. Additionally, an acoustic signal will sound at the test head.
• The pore detection sensitivity is adjustable. Depending on the setting, pores are indicated at short 20 to 50% voltage drops.
• Battery or line operation (switchable).
• Connector for external 12 V voltage supply.
• Connector for external On/Off switch of test voltage and isolated relay contacts to control acoustic signals or pulse counters.

4.11.4 Conductivity measurement

Electrical conductivity measurement of non-ferrous metals using SIGMASCOPE® SMP10

The electrical conductivity is an important material property that not only informs about how well a metal conducts electrical current but also provides information about its composition, microstructure or mechanical properties.

The electrical conductivity is measured using the eddy current method according to DIN EN 2004-1 and ASTM E 1004. The phase-sensitive measurement signal evaluation enables a contact-free determination of the electrical conductivity, for example, under paint or synthetic coatings of up to 500 µm in thickness. This also minimizes the influence of surface roughness.

Applications

• Measurement of the electrical conductivity of all non-magnetic metals, even stainless steel, EURO coins, etc.
• Measurement of the hardness and strength of heat-treated materials, e.g., aluminum alloys; inspection for heat damage.
• Measurement of the phosphor content in copper.
• Monitoring of deposition processes, e.g., for Cu-Cr-alloys.
• Determination of the degree of purity.
• Verifying the homogeneity of alloys.
• Scrap metal sorting.

Calibration standards

A high-precision measurement is required to determine the electrical conductivity. Accurate standards are required to calibrate the instrument because the measurement is a comparison using the eddy current method. These standards are available in certified versions for the entire conductivity range. Special standards are also available for testing EURO coins, for example.

The eddy currents generated by the magnetic field of the probe and utilized as the measurement effect are influenced by the electrical conductivity.

The penetration depth _ of the eddy currents is established by the measurement frequency f, which determines the minimum permissible thickness of the specimen.

The electrical conductivity measurement is employed for the production, processing or inspection of finished goods (e.g. EURO coins) made of NF metals. The electrical conductivity measurement is an important quality assurance component in the manufacture, maintenance or repair of airplanes.

4.11.4 Conductivity/resistivity sensors from METTLER TOLEDO

Conductivity measurements are carried out in industrial processes primarily to obtain information on total ionic concentrations (e.g. dissolved compounds) in aqueous solutions. Widely used applications are water purification, CIP/SIP control and the measurement of concentration levels in solutions. The measuring system consists of an appropriate inline sensor directly inserted or in a housing, a cable connected to a transmitter converting the received signals to a measurement result or forwarding it to a DCS.

InPro7001-VP (2-electrode sensor, for pure and ultrapure water applications)

The 2-electrode sensor InPro7001-VP is a highly accurate conductivity cell for pure and ultrapure water applications such as demineralization or desalination. The 12 mm body and the common PG 13.5 thread allows installation through most standard ports. The cell is steam-sterilizable up to 131°C.

Specifications

Cell constant 0.1 1/cm (nominal)

Temperature range -10…100 °C (14…212 °F)

Max. pressure (bar) 17 bar (25°C), 7 bar (95°C)

Max. pressure (psi) 250 psi (77 °F), 100 psi (203 °F)

Electrode material SS 316L/1.4435

Shaft material SS 316L/1.4435

O-rings, insulation Viton® (FDA), PEEK (FDA)

Accuracy (cell constant) ±1.0 %

Temperature accuracy ±0.25 °C at 25 °C

Temperature sensor Pt1000

Objectives

When you have completed this chapter you should be able to:

• Describe the Redox reaction
• Show how Redox measurement is performed
• List the main applications
• Demonstrate how calibration and checking procedures are carried out
• Explain the reasons why Redox measurement is not often applied

5.1 Introduction

Redox reaction is a shortened form of reduction-oxidation reaction. While originally applied to the increase or decrease in oxygen content, oxidation now refers to the loss of any electrons during a reaction and reduction refers to the gain of electrons. Oxidation and reduction always occur together — with one substance accepting the electrons that another loses. Thus in the reaction of sodium chloride, sodium loses an electron and is therefore oxidised whilst the chlorine atom gains an electron and is said to be reduced (Figure 5.1).

A measure of the power of a substance to gain electrons is called the Redox potential — sometimes referred to as ORP (Oxidation-Reduction Potential). A reducing agent that readily loses electrons will have a negative potential whilst oxidising agents will have a positive potential.

5.2 Redox measurement

Redox measurement is similar to pH measurement but with the pH-sensitive glass replaced by a metal electrode — usually gold or platinum (Figure 5.2 and 5.3).

The measured value is again determined by the Nernst equation, but with the difference that the variable parameter is not the pH value but the ratio of oxidising and reducing ions.

In practice, however, the Redox value measured is not determined exclusively by the Nernst equation since the surface quality of the metal used, pre-treatment and history of the metal electrode also influence the measurement — as do the constituents of the liquid.

Further, unlike pH and ion-selective electrodes, the response of a Redox electrode is non-specific. Thus, it is difficult to correlate the output with the concentration of a species, e.g. through calibration. For this reason, it is usual to use Redox potential measurements (in millivolts) as they stand to control a Redox process, rather than trying to convert them into concentration or other units.

It should be noted that the millivolt signal from a Redox pair also depend on the reference electrode used. A Redox potential measurement obtained using a calomel reference, for example, will differ by about 46 mV from that using a silver/silver chloride reference.

5.3 Applications

The measurement of Redox Potential is an important tool in the control of a variety of processes involving oxidation/ reduction reactions. The main applications include:

• Monitoring of chlorine proportioning in swimming pools (oxidising potential)
• Nitrite detoxication
• Cyanide detoxication
• Chromate detoxication

The conditions present in all four of these applications are straightforward for the Redox measurement. The metals selected for Redox measurement depend on the medium to be measured and are shown in Table 5.1.

The process of cyanide detoxication involves treatment with chlorine under alkaline conditions to oxidise the cyanides (CN^–) to cyanogen chloride (CNCL) which subsequently reacts to form cynate (CNC^–) and chloride (CL^–) ions. Both Redox and pH measurements are needed to optimise control of the process. The Redox measurement is used to ensure an adequate concentration of the oxidant, i.e. chlorine, is maintained (Figure 5.4).

In the example of the reduction of chromium wastes, the process involves reducing dichromate (Cr⁶⁺) to chromate (Cr³⁺) ions using sulphur dioxide (SO₂). The chromate is then precipitated as hydrated chromic oxide by increasing the pH and consequently removed as sludge. Although, in this case, the Redox potential measurement is used to control a reduction process, the essential features of the system remain very much the same (Figure 5.5).

5.4 Calibration/checking procedure

Since the measurement of Redox is a mV potential, chemical calibration is not required. However, solutions containing quinhydrone are often used to check the performance of the Redox pair. Quinhydrone is a compound of benzoquinone (Q) and hydroquinone (H₂Q) and forms a reversible reduction/oxidation system with a hydrogen ion:

Q + 2H⁺ + 2e^– ↔ H₂Q

Because the Redox system is pH dependent, this provides a convenient method of checking a Redox electrode pair. Due to oxidation, the system is unsuitable above pH 8.

5.5 Troubleshooting

5.6 Conclusion

Although the measurement of Redox potential is, in many cases, of great importance it is not, generally, a popular measurement unless applied on well tried and tested applications. The reasons include:

• Inadequate theoretical knowledge
• Difficulty in interpreting the results
• Practical difficulties in measurement
• Multiplicity of units (Redox potential; pe value; rH value).

8.6 Electrode maintenance

Sluggish response and low dissolved oxygen values are usually indicative of the need for service due to depletion of the electrolyte, fouling of the membrane or chemical deposition on the anode.

The sensor should be stripped, according to manufacturer’s instructions, and cleaned. When cleaning the sensor, copious flushing with water is necessary to remove all traces of foreign chemical compounds that might impair the sensor response.

The electrodes should never be scraped or scratched with mechanical devices since this will change their sensitivity and reproducibility. The gold cathode can be polished using a fine grade of alumina powder, as recommended by the manufacturer, using a soft polishing cloth. Following polishing, all traces of the alumina powder must be removed by applying a high velocity jet of water (e.g. from a syringe) to all sensor internal surfaces.

Once the sensor has been cleaned correctly, it is reassembled with fresh electrolyte and a new membrane.

The internal electrolyte is usually a 1.0 or 0.5 Molar solution of potassium hydroxide and potassium chloride or bromide. During the refilling procedure care must be taken to exclude any air bubbles. Such bubbles will affect the response time since the oxygen they contain will have to be consumed prior to measurement taking place. Furthermore, during periods of non-operation, oxygen will pass through the membrane and replenish the bubbles.

Different thicknesses of membrane are used for different concentrations of dissolved oxygen — with thicker membranes being used for higher concentrations. Theoretically, membranes from different manufacturers should be interchangeable. However, correct results may not be obtained since different diffusion potentials may be set up.

Following sensor refurbishment, a period of time in operation is required before calibration can take place, especially where a Faraday cell is used. This ensures the depletion of any oxygen dissolved in the electrolyte. A well-cleaned sensor should require not more than 1 hour to stabilise.

8.6.1 Sensor storage

When sensors are not in use, they should be immersed in freshly prepared 10% sodium sulphite solution in an air-tight container. This is in effect a zero-oxygen solution. When a sensor will not be in use for some time it is best to remove the internal electrolyte completely, rinse and store in demineralised water until required.

8.7 Applications

8.7.1 Wastewater applications

Control of wastewater is increasingly becoming important to avoid potential environmental and increasing legal consequences. Accurate pH, O₂, conductivity and turbidity measurements are critical in modern wastewater treatment. METTLER TOLEDO offers reliable, cost-effective solutions to control industrial effluents.

Accurate pH and conductivity measurement is important for determining the state of the raw effluent, for controlling the dosage of chemicals in neutralization and other necessary chemical treatment steps, as well as for monitoring the quality of the final discharge. O2 measuring systems ensure reliable concentration measurement in the biological treatment of industrial wastewater even under extreme conditions. Wastewater treatment procedures demand also measurement of suspended solids and turbidity is a key parameter for monitoring various process stages

8.7.2 Biotech and hygienic processes

METTLER TOLEDO/INGOLD offers a wide range of hygienic and sterilizable/auto-clavable measurement solutions for biotech and hygienic processes. Based on field experiences and expertises in hygienic design for CIP (Cleaning-in-Place) and SIP (Sterilization-in-Place) applications the solutions comply with the most stringent industry requirements. Critical biotech/hygienic applications in pharmaceutical, food and beverage are controlled by measuring pH, conductivity, dissolved oxygen, CO₂ and turbidity.

In biotechnology applications, the solutions ensure an optimum yield in the fermentation process, enable to control growth and metabolism of micro-organisms and cells and allow monitoring of the purification process with USP 24 compliant sensors. In hygienic applications in food & beverage EHEDG and 3-A approved sensors are for example used to control process oxygen level in breweries to avoid oxidation of substances leading to a change in taste and diminished shelf-life of beer.

8.7.4 Wastewater DO Sensors

InPro6050

InPro6050 sensors are designed for simultaneous measurement of dissolved oxygen and temperature values in the field of water applications. They are particularly suitable for applications in: aeration basins, outflow channels, denitrification tanks, environmental control, stations and fish farming tanks. The InPro6050 is a reliable O₂ sensor with built-in temperature device (RTD). The new sensor features an optimal price/performance ratio and allows continuous recording of dissolved oxygen content in all types of water applications.

Features and benefits

• Minimum calibration requirements — Easy air calibration before installing into process.
• Low maintenance requirements — Faster than ever before! The cleverly designed membrane module cuts full service time to just five minutes.
• Long sensor lifetime — TheTeflon-coated membrane prevents fouling, even under extremely dirty operating conditions.

Description

Together with the O2 4050e transmitter it provides reliable DO measurements during the biological water treatment or in fish farming applications. The robust, detachable VP (VarioPin) connector offers the hermetic advantages of fixed cable electrodes (IP68 watertightness), but avoids the need to interfere with the existing cable installation when disconnecting the electrode.

Specifications

Short description	simultaneous measurement of DO and temperature
Response time at 25 °C (air –> N2)	greater than 90 sec
Material in contact with medium	PPS, Silicone, Viton®, Teflon
Connector	VP connector (IP68 watertightness)
Pressure resistance measurement	2 bar (60 °C) / 29 psi (140 °F)
ISM	no

InPro6800

METTLER TOLEDO’s «Advanced Line» dissolved oxygen sensors are specifically designed for reliable in-line measurement in processes under sterile, hygienic conditions. InPro6800 sensors have been developed and manufactured in accordance with the most precise surface treatment standards to comply fully with EHEDG and FDA flat-surface recommendations for extremely high-level hygiene applications.

Features and benefits

• Reduced service time—Thanks to the easy-to-replace membrane body and quick-disconnect interior body.
• Long sensor life time—Durable and rugged sensor design for increased resistance to harsh environments.
• High process safety—FDA compliant materials of construction and easy-to-clean high-polished surface (N5 grade) to satisfy increasing regulatory requirements.
• Complies with the strictest regulations—Suitable for hygienic applications: EHEDG certified and 3-A compliant.
• Better process control—High accuracy and low detection limit of 6 ppb.
• Installation in hazardous zones—ATEX certificate: II 1/2GD EEx ia IIC T6/T5/T4/T3, IP6X T 69 °C / T 81 °C / T 109 °C / T 161 °C, SNCH 01 ATEX 3277X; FM approval: IS / I, II, III /1/ ABCDEFG / T6 Ta= 60 °C – 53 800 002; Entity.

Description

Modular Design: Saves time and money!

• Replacement of the anode/cathode assembly has been made simple with METTLER TOLEDO’s state-of-the-art interior sensor design, thereby reducing maintenance costs. No complicated service procedures: the quick-disconnect interior body can be replaced in less than 5 minutes! The fail-safe interior body design guarantees a good connection every time.

High process reliability guaranteed!

• Durable materials, robust design and easy replacement. The exclusive Ingold membrane was designed and developed by METTLER TOLEDO and has become the standard for quality throughout the industry.

No chance of contamination!

• The new O₂ sensors have been tested by leading standards institutes and definitively passed every requirement. METTLER TOLEDO sensors are the first O₂ sensors in the world to be awarded EHEDG (European Hygienic Equipment Design Group) certification and meet 3-A hygienic requirements.

Compatibility!

• The InPro6800 is fully compatible with the O₂ 4100e «Advanced Line» transmitter. The METTLER TOLEDO «Advanced Line» transmitters offers a wide range of performance benefits. Easy operation, robust construction and all needed functions make the advanced line an ideal choice for your demanding applications.

Specifications

Short description	specially intended for use in sterile/hygienic processes
Accuracy	1 % or 4 ppb
Sterilizable	yes
Autoclavable	yes
Material in contact with medium	stainless steel (AISI316L) with 3.1B certificates
Response time at 25 °C (air –> N2)	90 sec.
Mechanical pressure	max. 12 bar (174 psi) absolute
resistance
Detection limit	6 ppb
ISM	yes
ATEX Certification	II 1/2GD EEx ia IIC T6/T5/T4/T3, IP6X T 69 °C / T 81 °C / T 109 °C / T 161 °C, SNCH 01 ATEX 3277X
FM approval	IS / I, II, III /1/ ABCDEFG / T6 Ta=60 °C – 53 800 002; Entity

Objectives

When you have completed this chapter you should be able to:

• Explain the basic principles of chlorine chemistry
• Describe the measurement of total free chlorine and differentiate between the methods used
• Demonstrate the methods used for calibration
• Understand the application of total free chlorine measurement

9.1 Introduction

Chlorine (CL), a poisonous, yellowish-green gas, is an extremely reactive element that combines directly with most other elements and readily dissolves in water. In an aqueous solution, chlorine is a powerful oxidising agent and is used extensively in the control of sewage effluent; in the disinfection, taste and odour control of potable water; and in the disinfection and bacterial control of swimming pools.

Chlorine is also used in many industrial applications for bleaching, as a powerful oxidising agent in various manufacturing processes, in the control of unwanted algae and bacteria in cooling waters, and in the control of marine life in sea-cooled power stations.

However, chlorine is effective only at predetermined concentrations and under-dosing could prove to be ineffectual. Conversely, overdosing imparts an unpleasant flavor and odor to water. As a result, the levels must be carefully controlled.

9.2 Basic chlorine chemistry

When chlorine is dissolved in water, it may exist in one or more of three different forms —depending on the pH of the sample:

Figure 9.1 shows the equilibrium concentrations of each species with respect to pH.

In terms of disinfection, the form in which chlorine exists is extremely significant. Although hypochlorite (OCL^–) has a bactericide effect, it is 80 – 100 times less powerful than hypochlorous acid (HOCL). This means that for effective disinfection, it takes a concentration of 8 ppm of OCI^– to do the same job as a concentration of 0.1 ppm of HOCL.

As indicated in Figure 9.1 the equilibrium between hypochlorous acid (HOCL) and hypochlorite ion (OCL^–) is pH dependent — with the percentage of HOCL changing from 100% to 0% as the pH rises from 4.0 to 11.0. Figure 9.2 illustrates how this equilibrium changes with pH at 0°C and 20°C.

Although it would be desirable to be able to differentiate between these two forms of free chlorine, in practice the oxidising power of HOCL and OCL^– are similar and the chemical and electrochemical methods used to determine chlorine concentration cannot distinguish between them. Consequently, it has become traditional to express the residual chlorine, existing in water as hypochlorous acid and hypochlorite ions, as the ‘Total Free Chlorine’ or the ‘Free Available Residual Chlorine’ where:

Total Free Chlorine (TFC) = (OCL^–) + HOCL

Thus, the two forms of ‘Free Available Residual Chlorine’ exist together, with their relative concentrations depending on the pH of the solution and not on whether chlorine gas or bleach was added.

Table 9.1, which is derived directly from Figure 9.2, shows that if the pH of the water is 9.0, a free chlorine residual of 1 ppm contains only 0.029 ppm HOCL at 20°C. However, by adjusting the pH to 7.5 the HOCL concentration changes to 0.49 — a 17 fold increase in active disinfectant.

As a result of the foregoing, it is customary to run higher free residual chlorine levels with increasing pH, to ensure an HOCL level of 1 mg/ℓ (1 ppm) at all times in the water. Another way of expressing the disinfecting properties of chlorine is shown in Table 9.2 which illustrates the required free residual chlorine levels needed to achieve this level of hypochlorous acid.

Clearly, proper chlorine dosage should be combined with pH control in order to achieve the most efficient disinfection.

9.3 Measuring systems

The on-line measurement of chlorine may be accomplished either chemically or by amperometry. Two chemical methods are available: the colorimetric method and the titration method (see Chapter 10). The common colorimetric method relies on the color change of an ’indicator’ — diethyl-p-penylene-diamine (DPD) — to pink in the presence of chlorine and its compounds, with the color intensity related to the concentration of chlorine.

In the titration method, the color change produced by the indicator is titrated to colorless using ferrous ammonium sulphate. Both these methods require pH adjustment to between 6.2 and 6.5 and a chemical buffer is added to achieve this condition.

Since the determination of chlorine is pH dependent and the DPD method adjusts the sample pH for correct measurement to take place, some difficulties in the on-line determination of chlorine are experienced.

As a result, the most frequently used method for the on-line determination of chlorine is amperometry.

9.4 Measuring principle

Figure 9.3 illustrates a typical amperometric measurement cell which is similar, in many respects, to that used for dissolved oxygen.

The membrane-covered sensor consists of a gold cathode serving as the working electrode and a silver/silver chloride anode acting as the counter-electrode. Both electrodes are immersed in an electrolyte and a constant polarisation voltage of about 700 to 800 mV is applied between the anode and the cathode. At this voltage, only sufficiently strong oxidising agents, such as HOCL and OCL will pass through the membrane and react at the gold cathode.

Chlorine measurement is typically performed in media with a pH value in the range of 6.5 to 8.5. When the sensor is immersed in chlorinated water, the HCLO in the water diffuses through the membrane and reacts at the gold cathode:

HCLO + 2e^– → OH^– + CL^–

Silver oxidises to silver chloride at the anode to give:

2 Ag + 2 CL^– → 2 AgCL + 2e^–

The resulting current flow between the electrodes is directly proportional to the amount of Total Free Chlorine in the sample solution.

Because the chlorine molecules are consumed in the electrochemical reaction, they must be replaced by new ones in the process flow. As a result, the sensor requires a continuous flow of fresh sample at the sensor tip — the minimum recommended flow being 0.3 m/s second over the sensor membrane.

In one form of analyser, the measurement is made virtually independent of the flowrate by rotating the electrode at high speed.

9.5 Calibration

Polarisation at the cathode starts immediately after connecting the sensor to the measuring instrument and switching on the power. A period of approximately 30 to 60 minutes is required for polarisation to be completed.

Zeroing is normally accomplished with a water flow having zero chlorine content. This is achieved by placing an activated carbon filter in-line with the analyser.

Slope calibration is performed by passing chlorinated measuring water through the flow assembly and then determining the free effective chlorine concentration colorimetrically in the laboratory after sampling according to the DPD method.

9.6 Application

Chlorine dosing is preferably controlled by measuring the free chlorine content — with a pH value measuring and control system used to maintain the pH value constant. The Redox potential is monitored by a third measuring system, preferably equipped with an alarm limit contact. All three measured values — chlorine content, Redox potential and pH value — are continuously measured and recorded along with the water temperature.

9.6.1 Total chlorine measurement for drinking water disinfectant

In order to insure the supply of clean safe drinking water that meets the Safe Drinking Water Act (SDWA), and to provide a lasting measurable disinfectant residual in the water distribution system to protect against contamination, excess chlorine is added. The addition of chlorine beyond the point at which all chlorine demand is consumed, called the “breakpoint,” produces a free chlorine residual. Endress+Hauser’s chlorine systems are designed to measure this free chlorine residual. Water systems with extended distribution systems may add ammonia, either concurrently with, prior to, or following chlorine addition, to form chloramines. There are three forms of chloramines:

Monochloramine NH₃ + Cl₂ = NH₂Cl+HCl

Dichloramine NH₃ + 2Cl₂ = NHCl₂+2HCl

Trichloramine NH₃ + 3Cl₂ = NCl₃ + 3HCl

The formation of Monochloramine (NH2Cl), that is preferred, occurs at pH of 8.5 or higher. The formation of Dichloramine (NHCl₂) and Trichloramine (NCl₃) occurs at pH of 5.0 and lower. Chloramines, though not as potent a disinfectant as free chlorine, are more persistent in long distribution systems or transmission mains. Chloramine residuals, referred to as combined chlorine residuals plus any free chlorine residual present, constitute the total residual chlorine present in the water for disinfection. Endress+Hauser’s Total Chlorine System – CCS120 is designed to measure this total chlorine residual

9.7 Troubleshooting

Objectives

When you have completed this chapter you should be able to:

• Explain the basic principles of colorimetry
• Describe the basic principle of on-line titration

10.1 Introduction

Both colorimetry and titration have been used as standard laboratory analysis techniques for many years. However, modern computer and microprocessor control has brought these technologies into the area of on-line measurement to provide control over a variety of processes and applications that were, at one time, considered too difficult.

10.2 Colorimetry

In colorimetry one or more reagent solutions are added to the sample that converts the chemical substance being measured into a chemical complex of a predetermined color. The intensity of the color, which is measured optically, is proportional to the concentration of the chemical substance of interest.

The basic principle of the colorimetric method is shown in Figure 10.1. Once the sample has reacted with the reagent, the colored solution flows through a glass-sided measuring cell (cuvette) that allows light from a suitable light source to pass through the solution and onto a light detector. By virtue of the light absorbance by the solution, the output from the photocell is related to the intensity of the colored complex.

Because the chemical reaction with the sample is carefully chosen to produce a characteristic color, a particular band of wavelengths will be absorbed when white light is passed through the sample in the measuring cuvette.

To ensure that the measurement is reproducible, and to remove interference from colored substances in the sample, an optical filter, having a spectral response characteristic opposite to the solution, is placed in the light path (Figure 10.2).

On-line analysers use fixed optical filters that are available in two basic forms:

• Broad band filters — usually made of two or more layers of glass separated by a gel containing colored dyes.
• Interference filters — narrow band filters consisting of two layers of glass on whose inner surfaces a thin, semi-transparent metal film is deposited with an inner layer of a transparent material such as quartz or calcium fluoride.

At low concentration levels, the relationship between the concentration of a colored solution and light absorption is linear. At high absorbance values, however, there will be deviation from this relationship, as described by the Beer-Lambert Law, as shown in Figure 10.3. And as the concentration increases even further, a point is reached where the color generated absorbs virtually all the light and the response of the measurement reaches saturation.

Since the absorbance of the reacted solution is proportional to the path length of the cuvette multiplied by the concentration, the instrument range is dependent on: the path length of the cuvette; the point at which color saturation is reached; and the sensitivity, stability and dynamic range of the optical system and associated electronics.

A complete schematic of a typical colorimetric system is shown in Figure 10.4.

One of the major drawbacks of the colorimetric method for on-line measurement is that often the reactions are not instantaneous and require waiting periods to develop the complex. These waiting periods are achieved using delay coils. Reaction rates are also temperature dependant, so that fixed temperature conditions are necessary.

Despite this drawback, the colorimetric method has become the standard technique for the measurement of silica and phosphate. Both measurements use a chemical solution based on molybdenum, which will react with either silica or phosphate in the sample being measured, to form a yellow complex whose intensity is proportional to concentration.

The on-line measurement of chlorine can also be performed using the colorimetric method based on the color change of the indicator diethyl-p-penylene-diamine (DPD) to pink in the presence of chlorine and its compounds.

Since the determination of chlorine is pH dependent and the DPD method adjusts the sample pH in order for correct measurement to take place, some difficulties in the on-line determination of chlorine are experienced. Further disadvantages are that the analyser is complicated, has moving parts and consumes reagents which, in the case of the DPD, are costly. Furthermore, without knowledge of the pH of the sample the DPD method gives no indication of correct sterilisation.

10.3 Colorimetry applications

10.3.4 Spectrally tunable sources: colorimetry application

Colorimetric quantities of a light source can be measured with a colorimeter that has three or four channels consisting of detectors with spectral filters; the channels are designed to have relative spectral responsivities (RSRs) that mimic the color matching functions defined by the International Committee on Illumination (CIE). No colorimeter channel exactly matches each of the color matching functions; an example of the RSRs of colorimeter channels and the CIE-defined color matching functions is given in Figure 10.5. Due to this imperfect matching of the spectral responsivities, measurement errors are inevitable. These measurement errors can increase dramatically when the relative spectral power distribution (SPD) of a test source is dissimilar to that of the calibration source. Colorimetric quantities can also be measured with a spectroradiometer, with colorimetric quantities calculated from the measured spectral power distribution. Spectroradiometers theoretically do not have the spectral mismatch problem, but they are still susceptible to measurement errors due to wavelength error, stray light, and finite measurement bandwidth. These radiometric measurement errors introduce errors in calculated photometric and colorimetric quantities in a similar way that the spectral mismatch errors do in colorimeters (Figure 10.5).

10.4 Colorimetry troubleshooting

The Colorimeter requires no maintenance other than cleaning the parts that are active during color measurement, i.e. the Sapphire table, the fiber optic sensing input and, occasionally, the optical sphere, and outer housing.

10.5 Titration

In its simplest form, titration is a procedure in which the unknown concentration of a solution of a base is determined by allowing it to react with a known concentration of a solution of acid, or vice versa.

In the procedure, illustrated in Figure 10.6, a known volume of a solution of the base is dispensed into a flask from a burette (a) to which a few drops of a suitable indicator are then added (b). An acid solution of known concentration is placed into a burette (c) and slowly dispensed into the flask until the indicator just changes color (d) Usually this color transition, which marks the equivalence point (the point where the reactants and products are present in equivalent amounts), takes place with the addition of only one last drop of acid. Table 10.1 lists some other common acid-base indicators.

Although the example given above relates to the determination of the concentration of a base, titration may be applied to a variety of analyses in order to determine the volume of one solution required to react with a given solution of another solution.

Furthermore, the equivalence point may be determined, not just by a change in color but also by the sample undergoing some other change in state: e.g. cessation in precipitation (Ag⁺ plus CL^–); neutral pH as measured by a pH electrode; or by the use of ion selective electrodes.

Whilst for many years titration has remained a purely laboratory analysis, several companies have introduced automated titration systems for on-line analysis.

Such a system, illustrated in Figure 10.7, may be customised to perform a host of specific concentration analyses including: sulphide, nitrates, ammonia, fluoride, calcium, and cyanide.

Such a system may be used for both normal and abnormal titrations: e.g. titrations which are unstable, irreproducible and subject to interference such as pH-dependent ORP titrations. In such cases the analyser may be programmed to determine the endpoint mathematically by calculating the point of inflexion on the titration curve at the end of every analysis cycle. This means that the system is capable of titrating to one or more endpoints without any knowledge of the actual location of these points.

The analysis cycle starts by opening the drain valve (V3) and the rinse valve (V1) — allowing rinse water to simultaneously clean and drain the reactor for a programmed length of time. Following the closing of the rinse valve (V1), the sample valve (V2) opens, while valve (V3) stays open a few seconds longer in order to flush any remaining rinse-water droplets with fresh sample solution. Valve (V3) now closes and the sample volume is accurately adjusted with the built-in siphon (S). The reactor vessel is now charged with a defined volume of sample solution for titration.

The titration pump is now activated and the titration starts with the addition of the titrant solution until the endpoint is reached. The microprocessor now calculates the sample concentration from a knowledge of the titrant volume. Thus, if:

V_s . C_s = V_r . C_r

where:

V_s = sample volume

C_s = sample concentration

V_r = titrant volume

C_r = titrant concentration

Then:

C_s = k . V_r since V_s and C_r = constant.

10.6 Titration applications

10.6.1 Application of isothermal titration calorimetry

Isothermal titration calorimetry (ITC) is used to measure interactions between endocytic proteins and peptides, or proteins and other ligands such as lipid heads groups

For ITC experiments, two binding partners (e.g. protein A and protein B, in the same buffer and at known concentrations) are placed in the injection syringe and the ITC cell. The ITC cells are kept at a tiny temperature difference compared to a reference cell filled with buffer.

When protein B is injected, the two proteins interact. If this interaction is exothermic, the ITC uses less energy to heat the cell; if the interaction is endothermic, the ITC uses more energy to heat the cell. Precise measurement of the energy required to maintain temperature of the cell during the course of several injections leads to the calculation of free energy, enthalpy, entropy, Kd, and stoichiometry of binding from one single experiment. Furthermore the reaction can be measured without immobilisation or labeling of the binding partners.

10.6.3 Acid-base titration

In a typical acid-base titration experiment, the solution containing the analyte (an acid of unknown identity and/or concentration) is placed into a container, and the titrant (a base of accurately-known concentration) is slowly added from the burette in small increments (see Figure 10.9).

After the addition of each increment of base, the volume of base is carefully read from the burette (measured to the nearest hundredth of a milliliter) and the hydrogen ion (H⁺) concentration of the solution is measured with a pH meter. The pH meter gives the hydrogen ion concentration in terms of pH, which is simply the negative logarithm (common log, i.e. base 10) of the hydrogen ion concentration:

pH = -log[H⁺],

where [H⁺] represents the concentration of H⁺.

The hydrogen ion concentration is directly related to the amount of acid present in the solution at any particular step in the titration according to the following chemical reaction:

HA (an acid) H⁺ hydrogen ion + A^– (anion of the acid)

Chapter 1: Basic Chemistry

1. Describe an element.
2. Describe a compound.
3. Describe a mixture.
4. What is the difference between organic and inorganic chemistry?
5. What determines the atomic number of an atom?
6. What is the mass number of an atom?
7. Describe an isotope.
8. Describe the valence shell and the valence electron.
9. What are the main characteristics of metals?
10. What are the main characteristics of non-metals?
11. What is another term for semiconductor?
12. How many valence electrons does an alkaline metal have?
13. How many valence electrons does an alkaline earth metal have?
14. What are the characteristics of a halogen?
15. What is the difference between a cation and an anion?
16. Describe the difference between ionic bonding and covalent bonding.
17. What is the difference between molecular formula and empirical formula?
18. Describe stereo chemical formula.
19. Which suffix is used to describe a compound of metal and non-metallic elements?
20. What is the value of the atomic mass unit for oxygen?
21. Give the definition of a mole.
22. What is the main characteristic of an acid?
23. What is the main characteristic of a base?
24. Describe the oxidation-reduction reaction.

Chapter 2: Electrochemical Cells

1. Describe absorption and desorption.
2. What is the main effect of using a dilute acidic solution in an electrochemical cell?
3. What is one of the major defects of a simple voltaic cell? What is this term called?
4. Describe the action of a Daniell cell.
5. What is the effect of the porous barrier in a Daniell cell?
6. Describe the action of a salt bridge.
7. What is the electrochemical series?

Chapter 3: pH Measurement

1. What is the term a pH derived from?
2. Does a liquid with a value of below 7 pH exhibit an acidic or an alkaline property?
3. Define the term molarity.
4. Describe the basic principle of pH measurement.
5. What is hydration?
6. Why is a reference electrode required?
7. Gives three requirements that an ideal electrolyte solution should fulfill.
8. What is the other name for the reference junction?
9. Describe the construction of the reference junction.
10. Describe the Nernst equation.
11. Where would you use an antimony electrode?
12. Name five disadvantages of the antimony electrode.
13. As the ionic strength of a solution increases does the activity coefficient increase or decrease?
14. When does alkaline error occur?
15. Name the characteristics of high temperature, standard and low resistance electrodes?

Chapter 4: Conductivity Measurement

1. Define conductivity.
2. Which three factors define the conductivity of an electrolyte?
3. Explain the concept of cell constant.
4. In most solutions does the conductivity increase or decrease with increasing temperature?
5. Explain the difference between ‘solute’ and ‘solvent’, components in high purity water.
6. What solution has been developed to overcome the problems associated with traditional cells?

Chapter 5: Redox Measurement

1. Define the term ‘Redox reaction’.
2. What does ORP stand for and what is it a measure of?
3. What is the essential difference between Redox measurement and a pH measurement?
4. Name three applications for Redox measurement.
5. Give three reasons why the measurement Redox potential is generally not a popular measurement?

Chapter 6: Turbidity Measurement

1. Define turbidity.
2. What is a Raleigh scattering?
3. What are the two principles used in turbidity measurement?
4. Describe the Jackson Candle turbidimeter.
5. Describe a modern on-line absorptiometer.
6. How are the effects of the presence of dissolved color minimised?
7. What is the essential difference between an absorptiometer and a nephelometer?
8. Describe a method of overcoming the problem of the presence of air bubbles.
9. Define the Jackson Turbidity Unit and the difference between an FTU and an FNU.
10. What are the main advantages of the FAU and FNU units over FTU and NTU units?

Chapter 7: Hygrometry

1. What is water vapor?
2. Define the term ‘saturated vapor pressure’.
3. Explain why water boils at 78°C at an altitude of 7000 m.
4. Define partial vapor pressure.
5. Define relative humidity and show why an RH figure must be given with a temperature.
6. Show the difference between expressing the water vapor content in ppm (volume) and ppm (weight).
7. Describe how the hair element hygrometer works.
8. Describe the basic principle of the wet and dry bulb thermometers.
9. Describe the basic principle of the sling psychrometer.
10. What are the two main principles used in capacitive type sensors?
11. Describe the Regnault hygrometer.
12. Describe the basic principle of the cooled mirror dewpoint meter.

Chapter 8: Dissolved Oxygen Measurement

1. Describe the phenomenon of hydration.
2. Explain why the partial pressure of oxygen in a solution, in equilibrium with air or the gas mixture above it, is the same as the partial pressure of oxygen in the air or gas mixture.
3. What is the difference between partial pressure and concentration?
4. Describe the method of measuring DO using the membrane electrochemical sensor.
5. Describe the amperometric measuring cell.
6. Describe the galvanic measuring cell.
7. What are the main problems encountered when measuring DO and how are they overcome?
8. Describe the main precautions used for maintaining the DO electrode.

Chapter 9: Total Free Chlorine Measurement

1. Describe the main characteristics of chlorine.
2. When dissolved in water, chlorine may exist in three forms. What are they?
3. Describe how the equilibrium concentrations of each form changes with pH.
4. What are the three methods that may be used to measure chlorine?
5. Describe the amperometric method of chlorine measurement.

Chapter 10: On-Line Colorimetry and Titration

1. What is colorimetry?
2. Describe the basic colorimetric method.
3. How is interference from the effect of colored samples dealt with?
4. What are the main disadvantages of the colorimetric method?
5. Describe the basic titration procedure.

Chapter 1 – Basic Chemistry

(Exercises 1, 2, 3 and 4 should be done on their own, one at a time. They are applicable to Chapter 1 of the manual. If time is a problem, Exercise 3 can either be left out, or else kept back till the end of the workshop. Exercise 3 is also well suited to a class exercise. The exercise is put up on the screen, and a group answer from the class will hold sway.

This is NOT a chemistry workshop! These exercises have been included more to stimulate an interest in chemical aspects, rather than to test chemical abilities, and should be seen as such! As most of the delegates on this workshop will be either Engineers / Technicians, these exercises will probably be foreign to many, but quite stimulating none-the-less.)

Exercise 1

• Run the main directory page of the Analytical Instrument exercise programme. It should look as follows:

• Go to the section, titled BASIC CHEMISTRY.
• Select EQUATIONS.
• Select WORD EQUATIONS.
• Do this exercise to the end (bearing in mind that if you need to escape, press FILE, and then MAIN MENU).

Exercise 2

• Run the main directory page of the Analytical Instrument exercise programme.
• Go to the section, titled BASIC CHEMISTRY.
• Select EQUATIONS.
• Select ELEMENTS & SYMBOLS.
• Do this exercise to the end (bearing in mind that if you need to escape, press FILE, and then MAIN MENU).

Exercise 3

• Run the main directory page of the Analytical Instrument exercise programme.
• Go to the section, titled BASIC CHEMISTRY.
• Select EQUATIONS.
• Select SYMBOL EQUATIONS.
• Do this exercise to the end (bearing in mind that if you need to escape, press FILE, and then MAIN MENU).

Exercise 4

• Run the main directory page of the Analytical Instrument exercise programme.
• Go to the section, titled BASIC CHEMISTRY.
• Select EQUATIONS.
• Go to the section, titled INTERACTIVE SIMULATIONS. (This might not appear obvious, immediately. But, look for the tab, near the centre top of the scrren.)
• Select EQUATION BALANCER.
• Play around, until you can get the following combination to work. By pressing the OK button, you will see a green check.

Your task is now to do the same, but using:

until you have a balanced equation.

Chapter 2 – Electrochemical Cells

Definitions

Exercise 5

(Exercise 5 pertains to Electrochemical Cells, and can be done on it’s own.)

• Run the main directory page of the Analytical Instrument exercise programme.
• Go to the section, titled ELECTROCHEMICAL CELLS.
• Select ELECTROCHEMICAL CELLS.
• If you need to escape, press FILE, and then MAIN MENU.
• Keep pressing on the grey DICE, until you come up with a combination of copper and zinc.
• Look very carefully at the voltages that have been displayed at the top. Now, we know that conventional current flows from positive to negative, and that electron flow is from negative to positive.
• Thus, common sense dictates that the electrons will flow from the zinc to the copper. Press the button, to confirm this. Press the same button to take this away.
• The voltage difference between the two electrodes should be the highest voltage minus the lowest voltage. Therefore, 0.34 – (-0.76) = 1.7V. Press the button, to confirm this. Press the same button to take this away.
• The positive side should, therefore be at the copper, and the negative side at the zinc. Press the button, to confirm this. Press the same button to take this away.
• Current flows from the anode to the cathode. Thus, zinc should be the anode, and copper the cathode. Press the button, to confirm this. Press the same button to take this away.
• Common sense dictates that oxidation will take place at the zinc, and reduction at the copper. Press the button, to confirm this. Press the same button to take this away.
• Finally, press the and buttons, to see what happens on both sides of the cell.
• Now, let’s see if you can get one right on your own.
• Keep pressing on the grey DICE, until you come up with a combination of zinc and platinum. (There will be different combinations of zinc and platinum, and various delegates may select different combinations. This is not serious if the experiment of one delegate looks different to that of another, as the methodology will remain exactly the same.)
• First, try to answer the questions below, without pressing any buttons.
  • In what direction will the electrons flow?
  • What is the cell voltage?
  • Distinguish between the positive and the negative terminals.
  • Identify the anode and the cathode.
  • What would the equations be, on both sides?
• Now, press the buttons to get the right answers.

Exercise 6

(Exercises 6 and 7 also pertain to Chapter 2, but can be done together, in quick succession.)

• Run the main directory page of the Analytical Instrument exercise programme.
• Go to the section, titled ELECTROCHEMICAL CELLS.
• Select VOLTAIC CELL (and, depending upon your OS, you may have to select the option to allow the content to be displayed).

You should see the following picture.

• On the left hand side, select copper (under METALS) and the copper solution (under SOLUTIONS). On the right hand side, select zinc (under METALS) and the zinc solution (under SOLUTIONS).
• Now, click on the ON / OFF switch (once), on the top right of the diagram, to see what happens. When the electrons eventually stop moving, you can switch the process off, and then start again, if you so desire. The whole aim of this exercise is to observe the movement of the electrons, and what actually happens, inside the solutions, themselves. Please do take specific notice of when the electrons stop flowing!

Exercise 7

• Shut the page down, and start again.
• Run the main directory page of the Analytical Instrument exercise programme.
• Go to the section, titled ELECTROCHEMICAL CELLS.
• Select VOLTAIC CELL (and, depending upon your OS, you may have to select the option to allow the content to be displayed).
• Select copper on the left hand side, and silver on the right hand side.
• Observe what happens, when the process is switched on.

Chapter 3 – pH Measurement

(Exercises 8, 9 & 10 pertain to the chapter on pH measurement, but can be done together, in quick succession.)

Definitions
Mole: The mole is the SI unit of “amount of substance”. This “amount of substance” is defined as that which contains the Avogadro constant of particles (atoms, ions, or molecules).

Avogadro constant: The Avogadro constant is the number of particles in 1 mole of substance. It has the value 6.025 x 10²³, as defined by the number of atoms in 12g of the carbon-12 isotope.

Molarity: A molar solution contains 1 mole of solute in 1 dm ³ (1 litre) of solution. It is an indication of the concentration of the solution. Molar solutions have units of mol dm ⁻³ of solution (1 mol dm ⁻³). Dilute bench acids or alkalis are around 2 mol dm ⁻³ and concentrated acids are around 10 mol dm ⁻³.

Exercise 8

• Run the main directory page of the Analytical Instrument exercise programme.
• Go to the section, titled pH Measurement.
• Select pH METER (and, depending upon your OS, you may have to select the option to allow the content to be displayed).

You should see the following picture.

• Select ACID, under SOLUTIONS.
• Select hydrochloric acid (HCl).
• Under MOLARITY, dial in 2 mol dm ⁻³, and make use of a volume of 90 ml.
• Press INSERT PROBES, and observe the pH.
• Press REMOVE PROBES.
• Under MOLARITY, dial in 10 mol dm ⁻³, and make use of a volume of 90 ml.
• Press INSERT PROBES, and observe the pH.
• Does this tie up with the statement mentioned above?

Exercise 9

Using exactly the same figure, please complete the following table:

Exercise 10

• Remove the probes.
• Select ACID.
• Select HNO3.
• What molarity is required (? mol dm ⁻³), to produce a pH of 1.37, in a volume of 120 ml?

(This does appear to be pretty much trial and error, and that is exactly what it is. In the real world, it would be more proper to do a formal calculation here. However, this would be the domain of the Chemical Engineer, and most certainly would not be expected from the Instrumentation person, who would not typically have access to the formulae!)

Chapter 4 – Conductivity Measurement

Definitions
Anion: Anions are atoms or molecules containing more electrons than protons, and so carrying a negative charge. An anion is a negatively charged ion that is attracted to the anode during electrolysis. All non-metal ions and most radicals (except the ammonia ion) are anions.

Cation: Cations are atoms or molecules containing fewer electrons than protons, and so carrying a positive charge. A cation is a positively charged ion that is attracted to the cathode during electrolysis. All metal ions and the hydrogen ion are cations.

Additional background
When a mild electric field (Volts, E) is applied across a pair of electrodes immersed in a solution, there is a tendency for current (Amperes, I) to flow. This current is carried through the metal wires by electrons, but electrons are not able to move through the solution. The current through the solution must be carried by ions. Part of the current is carried by negatively-charged anions moving in the same direction as the electrons. The rest of the current is carried by positively-charged cations moving in the opposite direction.

The amount of current that is observed depends on a number of factors:

1. The current (I) is directly proportional to the voltage (E).
2. The current (I) is directly proportional to the area of the electrodes (A).
3. The current (I) is inversely proportional to the distance between the electrodes (d).
4. The current (I) is roughly proportional to the concentrations of the ions (C).

This suggests an equation relating these terms: I = constant x E x A x C/d , in which the constant depends on the solvent, the nature of the ions, and the temperature.

The ratio of voltage to current is defined as the electrical resistance (E/I = R), so the ratio of current to voltage is the reciprocal of the resistance: I/E = 1/R = constant x C x (A/d) .

The quantity A/d is determined by the construction of the electrodes, and is called the cell constant (K_cell, cm) with dimensions of cm²/cm.

Conductivity is defined as 1/R. It has the units of ohms^-1, and was originally called mhos, but is now defined as Siemens (S).

The specific conductivity (K, kappa) is defined as the conductivity between electrodes of 1 cm² area and 1 cm apart: K = d/AR = 1/(R x K_cell) , with dimension S/cm , and K = constant x C .

When the concentration of the ions is given in moles/cm³, the constant is called the molar conductivity and has the dimension of S-cm²/mole.

It is more common to use a term called the equivalent conductivity (ƛ , lambda , S-cm²/equivalent)

K = ƛ x C x N , in which N is a simple factor for converting from moles to equivalents:

N = 1 for 1:1 electrolytes such as HCl, KBr, NaNO₃, etc.

N = 2 for H₂SO₄, BaCl₂, PbSO₄

N = 3 for Al(OH)₃, H₃PO₄

N = 6 for Al₂(SO₄)₃

Exercise 11

(Exercises 11, 12 and 13 are applicable to Conductivity. Exercises 11 and 12 could be done in quick succession, and question 13 separately.)

• Run the main directory page of the Analytical Instrument exercise programme.
• Go to the section, titled CONDUCTIVITY.
• Select EXPERIMENT (and, depending upon your OS, you may have to select the option to allow the content to be displayed).
• We will now select a cation and an anion, to form an ionic compound. For the sake of this experiment, select Sodium and Chloride. This experiment will test the conductivity at different concentrations.
• Click on the button, and type in a value of 0.00010 moles / L (which, by the way, would be 0.10 moles / cm³).
• Click on the dark blue bar of the instrument on the screen, and jot down the conductivity.
• Click on the button, and type in a value of 0.00100 moles / L (which, by the way, would be 1.00 moles / cm³).
• Click on the dark blue bar of the instrument on the screen, and jot down the conductivity.
• Click on the button, and type in a value of 0.01000 moles / L (which, by the way, would be 10.0 moles / cm³).
• Click on the dark blue bar of the instrument on the screen, and jot down the conductivity.

Exercise 12

• If the program is already open, then jump down to the 4^th step from here. If not, proceed downwards.
• Run the main directory page of the Analytical Instrument exercise programme.
• Go to the section, titled CONDUCTIVITY.
• Select EXPERIMENT (and, depending upon your OS, you may have to select the option to allow the content to be displayed).
• Select Silver Acetate.
• Click on the button, and type in a value of 0.0001000 moles / L. Determine it’s conductivity.
• Select Silver Nitrate.
• Click on the button, and type in a value of 0.001000 moles / L. Determine it’s conductivity.
• Select Silver Sulphate.
• Click on the button, and type in a value of 0.01000 moles / L. Determine it’s conductivity.

Exercise 13

Precise measurements show that the equivalent conductivity varies slightly with concentration, approaching a limiting value at infinite dilution (ƛ_o), and increases with increasing temperature. The concentration of ions is equal to the equivalent concentration for an ionic compounds which ionizes completely (strong electrolyte), but for a partially-ionized compound (weak electrolyte) the concentration of ions may be much smaller than the equivalent concentration of the compound.

• Run the main directory page of the Analytical Instrument exercise programme.
• Go to the section, titled CONDUCTIVITY.
• Select SIMULATION (and, depending upon your OS, you may have to select the option to allow the content to be displayed).
• Under SOLUTIONS, select ACID and then select H2SO4.
• Keep the volume at 20.0 mL. The molarity is 1.0 M.
• Press and observe the intensity of the light.
• Press
• Change the volume to 120.0 mL. The molarity is still at 1.0 M, so in effect, the strength has not changed (even though the volume has), and therefore the flow of current (due to the conductivity) should be identical.
• Press and observe the intensity of the light remains the same.

• Under SOLUTIONS, select ACID and then select HNO2.
• Keep the volume at 20.0 mL. The molarity is 1.0 M.
• Press and observe the intensity of the light.
• Press
• Change the volume to 120.0 mL. The molarity is still at 1.0 M, so in effect, the strength has not changed (even though the volume has), and therefore the flow of current (due to the conductivity) should be identical.
• Press and observe the intensity of the light remains the same (although it is probably not as bright, as was seen with the H2SO4).

• Please repeat the experiment, using HC3H5O3 (under Acid) and C6H5NH2 (under Base). Just as a matter of interest, their conductivities differ, and this is evident in the flow of current through the light.

Chapter 5 – REDOX Measurement

Definitions
Redox: Involves elements in the reaction gaining and losing electrons. (Rusting is a redox reaction: as iron loses electrons and oxygen gains them, iron (III) oxide is formed, which is rust.)

This particular experiment (exercise 14) also fits in well with a titration test.
Titration: Titration is a technique which can be used for the neutralisation of an acid and an alkali. The alkali is accurately measured, using a pipette and placed in a conical flask with an indicator. The acid is then slowly added from a burette. When the indicator changes colour, the neutral point has been reached. (This sequence could also be reversed.)

(Exercise 14 can be applicable to REDOX measurement, or On-Line Colorimetry and Titration, depending on the application.)

Exercise 14

• Run the main directory page of the Analytical Instrument exercise programme.
• Select REDOX MEASUREMENT (near the bottom right, and, depending upon your OS, you may have to select the option to allow the content to be displayed).
• The instructions for this experiment (as originally received by the compiler of the exercises) are neither complete, nor has the designer of the software shared his formulae with us. None-the-less, one can do a certain section of the experiment, just to get an appreciation of what actually gets done, during such an analysis.
• Press and select the option that has as the Oxidizing agent and as the reducing agent.
• Please note that this programme runs with a random molarity for (and, thus, delegates will not receive identical results).
• Have a look at how many ml of you need to add, to create a redox titration of the product. (You will find that the colour of the product in the beaker at the bottom will change from light blue to purple. When you start getting close, the colour in the bottom will start to change colour, and at this point it may be wise to use the button.)

Chapter 6 – Turbidity Measurement

Definition
Rayleigh Scattering: Rayleigh scattering (named after the British physicist Lord Rayleigh) is the elastic scattering of light or other electromagnetic radiation by particles much smaller than the wavelength of the light, which may be individual atoms or molecules. It can occur when light travels through transparent solids and liquids, but is most prominently seen in gases. Rayleigh scattering is a function of the electric polarizability of the particles.

Rayleigh scattering of sunlight in the atmosphere causes diffuse sky radiation, which is the reason for the blue color of the sky and the yellow tone of the sun itself.

Rayleigh scattering causes the blue hue of the daytime sky and the reddening of the sun at sunset.

Rayleigh scattering is more evident after sunset. This picture was taken about one hour after sunset at 500m altitude, looking at the horizon where the sun had set.

All of this is courtesy of https://en.wikipedia.org/wiki/Raleigh_scattering

Absorption spectroscopy (Absorptiometer)
This example discusses the general principle using visible light as specific example. A white beam source — emitting light of multiple wavelengths — is focused on a sample (the complementary color pairs are indicated by the yellow dotted lines. Upon striking the sample, photons that match the energy gap of the molecules present (green light in this example) is absorbed in order to excite the molecule. Other photons transmits unaffected and, if the radiation is in the visible region (400-700nm), the transmitted light appears as its complementary color. By comparing the attenuation of the transmitted light with the incident, an absorption spectra can be obtained.

Courtesy of https://en.wikipedia.org/wiki/Absorption_(electromagnetic_radiation)

Nephelometer
A nephelometer is a stationary or portable instrument for measuring suspended particulates in a liquid or gas colloid. A nephelometer measures suspended particulates by employing a light beam (source beam) and a light detector set to one side (often 90°) of the source beam. Particle density is then a function of the light reflected into the detector from the particles. To some extent, how much light reflects for a given density of particles is dependent upon properties of the particles such as their shape, color, and reflectivity.

Courtesy of https://en.wikipedia.org/wiki/Nephelometer

Chapter 7 – Hygrometry

(Exercise 15 is applicable to Hygrometry, and can be done on its own.)

Exercise 15

Immediately below, one can find a table which was produced by the American Navy, for military personnel. It gives a Wet Bulb Globe Temperature index (in degrees F), and refers to work that can be conducted. The work categories that have been described are i) Easy Work, ii) Moderate Work and iii) Hard Work. There is also reference to how much water is recommended be consumed, to keep hydration levels up

• Run the HTML file marked as “rH”, which will be provided to you, by your instructor.
• Imagine you are in charge of builders, based in New York, USA. The actual station pressure in New York would be 101.3 kPa, which would equate to 1013 millibars.
• Check out on the table how long you would allow the workers to perform moderate building work, as well as how long you would allow them to rest, under various conditions. (Hint: you would observe the “wet bulb” temperature in degrees Fahrenheit, and use this to check up on the provided table.)
  • The temperature is 28 degrees C, and the relative humidity at 65%.
  • The temperature is 29 degrees C, and the relative humidity at 65%.
  • The temperature is 30 degrees C, and the relative humidity at 70%.
  • The temperature is 32 degrees C, and the relative humidity at 80%.
  • The temperature is 33 degrees C, and the relative humidity at 80%.
  • The temperature is 33 degrees C, and the relative humidity at 85%.
  • The temperature is 33 degrees C, and the relative humidity at 88%.
  • The temperature is 33 degrees C, and the relative humidity at 95%.
• Should the building workers now be moved off to Dubai, in the United Arab Emirates, where the temperatures and humidity levels are more severe in the summer months, and hard building work is expected to be performed, please repeat the experiment for the following conditions, for a Hard Work:
  • The temperature is 42 degrees C, and the relative humidity at 40%.
  • The temperature is 42 degrees C, and the relative humidity at 50%.
  • The temperature is 42 degrees C, and the relative humidity at 60%.

Chapter 10 – On-Line Colorimetry and Titration

(Exercise 16 is applicable to On-Line Colorimetry and Titration, and can be done on its own.)

Exercise 16

Spectrophotometry is the quantitative measurement of the reflection or transmission properties of a material as a function of wavelength. It is more specific than the general term electromagnetic spectroscopy in that spectrophotometry deals with visible light, near-ultraviolet, and near-infrared, but does not cover time-resolved spectroscopic techniques.

Spectrophotometry involves the use of a spectrophotometer. A spectrophotometer is a photometer (a device for measuring light intensity) that can measure intensity as a function of the light source wavelength. Important features of spectrophotometers are spectral bandwidth and linear range of absorption or reflectance measurement.

https://en.wikipedia.org/wiki/Spectrophotometry

• Run the main directory page of the Analytical Instrument exercise programme.
• Go to the section, titled SPECTROMETRY.
• Depending upon your OS, you may have to select the option to allow the content to be displayed.
• It may show you a message with a code. Just press OK.
• Turn on the instrument, by clicking on the knob on the left front plate, facing you.
• Manipulate the yellow arrows, on this same knob, until the TRANS value shows 0.0
• Click on the red arrow, next to the red block, to get a sample. Keep clicking on the UP arrow UNTIL you get 100% of this sample.
• Click on the sample bottle (right), to put this sample of yours into Sample Tube 1.
• Click on the blue arrow, next to the blue block, to get a sample. Keep clicking on the UP arrow UNTIL you get 100% of this sample.
• Click on the sample bottle (right), to put this sample of yours into Sample Tube 2.
• Click on a combination of the red and blue, to make up your own sample. (Good advice – just for a reference, jot down the percentage combinations you will use. For example, you might use 40% of the red colour, and 70% of the blue colour – but this is entirely up to you. In my example (40 and 70), then 40 / (40 +70) * 100% = 36.36%, and 70 / (40 +70) * 100% = 63.65%. I will refer to these later, just to see if my instruments is accurate or not.)
• Click on the sample bottle (right), to put this sample of yours into Sample Tube 3.
• Click on Sample Tube 1, to place this into the spectrophotometer.
• Click on the SCAN button (not on the machine), and then close the data window that pops up, showing the values for Sample 1.
• Click where Sample Tube 1 is stored, to return it.
• Click on Sample Tube 2, to place this into the spectrophotometer.
• Click on the SCAN button (not on the machine), and then close the data window that pops up, showing the values for Sample 1 and Sample 2.
• Select a wave length for Sample 1, where the absorbance of the red dye is strong and that of the blue dye is negligible. (This is where the red value is going to be high, and the blue value is going to be as close to zero as possible.) Write this down.
• Select a wave length for Sample 2, where the absorbance of the blue dye is strong and that of the red dye is negligible. (This is where the blue value is going to be high, and the red value is going to be as close to zero as possible.) Write this down.
• Set the wavelength for what you have chosen for Sample 1.
• Click on the space of Sample 2, to remove it from the machine.
• Click on Sample 0 to place it into the machine.
• Operate the knob (front right) and use the yellow arrows to set this to 0.0.
• Place each of the tubes into the machine, press the WAVELENGTH button and put in your first wave length value, press SCAN, clear the log, and put the test tube back. After testing the third tube, you should be able to jot down the following:
  • Sample Tube 1 (the red one) = …………..
  • Sample Tube 2 (the blue one) = …………..
  • Sample Tube 3 (your one) = …………..
• Get the percentage of Tube 3 to Tube 1 (as a clarification, divide the value of Tube 3 by the value of Tube 1, and multiply the answer by 100), and this percentage of the red solution you used in your sample, you should be very close.
• Remove Tube 3.
• Set the wavelength for what you have chosen for Sample 2.
• Click on Sample 0 to place it into the machine.
• Operate the knob (front right) and use the yellow arrows to set this to 0.0.
• Place each of the tubes into the machine, press the WAVELENGTH button and put in your second wave length value, press SCAN, clear the log, and put the test tube back. After testing the third tube, you should be able to jot down the following:
  • Sample Tube 1 (the red one) = …………..
  • Sample Tube 2 (the blue one) = …………..
  • Sample Tube 3 (your one) = …………..

(Just as a matter of interest, you might have noticed that the instrument actually calculated the values for a whole host of wavelengths, and the repetition with the second wave length was actually not necessary. However, this is only unique to this particular computer model, and DOES NOT happen in real life!!! So, you did the full experiment, just for the sake of completeness.)

• Get the percentage of Tube 3 to Tube 2 (as a clarification, divide the value of Tube 3 by the value of Tube 2, and multiply the answer by 100), and this percentage of the blue solution you used in your sample, you should be very close.

Now, one could calculate the actual amount of the product in the solutions, but this would, once again, be out of the scope of the Instrument Engineer / Technician.

Exercise 1

Rust

Compound

Oxygen

Magnesium Oxide

Solid

Gas

Hydrogen(g)

Water(l)

Liquid

Sodium Chloride

Iron Sulphide

Sulphur Dioxide(g)

Sulphuric Acid

Carbon Dioxide

Water

Methane(g)

Oxygen(g)

Oxygen

Water

Oxygen

Carbon Dioxide(g)

Exercise 2

Oxygen

Atoms

Molecules

Calcium

Magnesium

Oxygen

Fluorine

Helium

Potassium

Radium

Phosphorus

Silicon

Copper – wires

Steel – Bridges

Magnesium – aircraft

Gold – Jewellery

Aluminium – pots

Tungsten – lights

N

Ca

Na

Carbon

Fluorine

Krypton

Potassium – K

Lead – Pb

Copper – Cu

Sodium – Na

Platinum – Pt

Phosphorus – P

Exercise 3

Hydrogen

H2

O

7

6

MgO

O2

O

2

2

Na(s) Cl2(g) NaCl(s)

Cl

2

2

4

Al(s) O2(g) Al2O3(s)

2

3

4

X2	X2Y2
4Fe	X3Y2
2Cu	X1Y2
2K		X2Y1
N2		X2Y3
FeO3	X3Y3

Exercise 4

Exercise 5

Exercise 6

Just an observation of the delegates.

Exercise 7

Just an observation of the delegates.

Exercise 8

2.69

1.99

Exercise 9

Exercise 10

42 mol dm ⁻³

Exercise 11

Approximately 11.9 μS

Approximately 111.8 μS

Approximately 1064 μS

Exercise 12

Approximately 9.8 μS

Approximately 91.3 μS

Approximately 862.5 μS

Exercise 13

Just an observation of the delegates.

Exercise 14

Just an observation of the delegates.

Exercise 15

73.35 Degrees Fahrenheit, and there would be no Limit on the work / rest conditions.

74.95 Degrees Fahrenheit, and there would be no Limit on the work / rest conditions.

78.03 Degrees Fahrenheit, and there would be no Limit on the work / rest conditions.

84.25 Degrees Fahrenheit, and there would be 50 minutes work followed by 10 minutes rest in each hour.

85.95 Degrees Fahrenheit, and there would be 40 minutes work followed by 20 minutes rest in each hour.

87.37 Degrees Fahrenheit, and there would be 40 minutes work followed by 20 minutes rest in each hour.

88.21 Degrees Fahrenheit, and there would be 30 minutes work followed by 30 minutes rest in each hour.

90.10 Degrees Fahrenheit, and there would be 20 minutes work followed by 40 minutes rest in each hour.

85.26 Degrees Fahrenheit, and there would be no 30 minutes work followed by 30 minutes rest in each hour.

89.85 Degrees Fahrenheit, and there would be 20 minutes work followed by 40 minutes rest in each hour.

93.99 Degrees Fahrenheit, and there would be 10 minutes work followed by 50 minutes rest in each hour.

Exercise 16

Just an observation of the delegates.

Objectives

When you have completed this chapter you should be able to understand about:

• Reasons for silica content in water
• Forms of silica presence in water
• The Method of Analysis
• Principle of Operation
• Industrial Silica Analyzer
• Industrial Applications
• Benefits
• Trouble shooting
• Control of Silica in feed water

1.1 Introduction

Silica analyzers in modern power plants alert operators to harmful silica concentrations in the water-steam turbine cycle. In the boiler, silica forms silicate deposits that interface with heat transfer and are difficult to remove. In the turbine, silica builds up on the blades and causes drastic decreases in efficiency. Power plants with high boilers closely monitor silica concentrations through silica analyzers to avoid these problems

Silica forms a dense porcelain-like scaling that cannot even be removed with acid. Silica scaling also have a very low thermal conductivity. Because of its low thermal conductivity, a 0.5 mm build up of silica can reduce thermal transfer by 28% , reducing efficiency , leading to hot spots and ultimately ruptures.

In order to assure optimum turbine performance, continuous monitoring of silica in superheated steam, boiler water and feed water is of utmost importance. The only way to effectively control silica build-up is through effective monitoring through Silica Analyzer.

1.2 Reasons for silica content in water

Among many contaminants in the steam/water circuit, silica plays an important role in process monitoring, mainly because it is highly soluble in steam and extremely difficult to remove from steam/water. Silica is a contaminant that appears in many potential external and internal entry points.

External contamination:

• Raw water ingress (demine plant) or drain line mix-up (wearing of seal or incorrect installation)
• Use of silicon based lubricants and oils (result from leaky seals in water system or turbine oil leaks or by silicon-based coatings on tubes used in replacement activity)
• Feed water system (e.g. Unreacted silicon) or chemical dosing or reagent problems ( e.g. Caustic addition after resin bed regeneration)

Internal contamination:

• The condenser dust (a build-up of paint, quartz and grease, overhaul materials such as gasket materials, silicone sealant and kitty litter are potential sources of condenser dust)
• Oil spill absorbent materials ( kitty litter, diatomaceous earth)
• In case of required replacement an open boiler tube could cause fly ash contamination or refractory material
• Blasting materials ( need to clean the LP turbine or scale from the process)
• Accidental misplacement of materials (caused by work practices and poor housekeeping)

1.3 Silica presence in water

Silica or silicon dioxide ( SiO2) is found in abundance in nature. Common forms of silica include sand and quartz. Silica occurs naturally at levels ranging from a few ppm to more than 200 ppm. It is one of the most prevalent elements in the world.

Silica (silicon dioxide) in some cases is an anion. The chemistry of silica is complex and somewhat unpredictable subject.

The “Total Silica” content of a water is composed of “Reactive Silica” and “Unreactive Silica”.

Reactive Silica (e.g. Silicates SiO4) is dissolved silica that is slightly ionized and has not been polymerized into a long chain. Reactive silica is the form that RO and ion exchange chemists hope for. Reactive silica is the form of silica to be used in RO projection programs. Reactive silica, though it has anionic characteristics, is not counted as an anion in terms of balancing a water analysis, but it is counted as part of total TDS( Total Dissolved Solids).

Unreactive silica is polymerized or colloidal silica, acting more like a solid than a dissolved ion. Silica, in the colloidal form can be removed by a RO but it can cause colloidal fouling of the front-end of RO. Colloidal silica, with sizes as small as 0.008 micron can be measured empirically by the SDI ( Slit Density Index) test, but only that portion that is larger than 0.45 micron or larger.

Particulate silica compounds (e.g. clays, slits and sand) are usually 1 micron or larger and can be measured using the SDI test.

Polymerized silica, which uses silicon dioxide as the building block, exists in nature (e.g. quartzes and agates).

Silica, in the polymerized form, also results from exceeding the reactive silica saturation level. The solubility of reactive silica is typically limited to 200-300% with the use of silica dispersant. Reactive silica solubility increases with increasing temperature, increases at pH less than 7.0 or more than 7.8 and decreases in the presence of ion which acts catalysts in the polymerization of silica. Silica rejection is pH sensitive, with increasing rejection at a more basic pH as the reactive silica exists more in the salt form than in the acidic form:

• Ionic Silica (reactive)
• (OH)3Si-O-
• Ionic or reactive silica exists in an SiO2 complex. It is not strongly-charged ion and therefore is not easily removed by ion exchange. However, when concentrated to levels above 100 pm, ionic silica may polymerize to form a colloid.
• Colloidal Silica (nonreactive)
• -SiO2-O-SiO2-O-SiO2-

At concentrations over 100 ppm, silica may form colloids of 20,000 molecular weight and larger, still too small to be effectively removed by a particle filter. Colloidal silica is easily removed with ultra filtration or can be reduced by chemical treatment (lime softening)

Colloidal silica can lower the efficiency of filtration systems such as reverse osmosis.

Any silica can affect yields in semi conductor manufacturing and is a major concern in a high pressure boiler systems.

Effect of silica in boiler feed water

Despite having no direct corrosive effect on plant, silica can form extremely hard and dense scales in the boiler and turbine, hampering heat transfer efficiency and increasing the risk of mechanical failure such as turbine blade malfunction. Silica entering a high pressure boiler can concentrate very quickly.

Just 1 ppm of silica in the feed water for a 500W boiler evaporating 1,500 tonnes of water per hour will result in 1 tonne of silica being deposited in the boiler in just one month.

As dissolved silica is only weakly ionized, it is difficult to detect by conductivity measurement. For this reason, dedicated silica analyzers are necessary if accurate information is to be obtained.

Depending on the type of power plant, typical sampling points for silica monitoring include the water treatment plant, the boiler drum and saturated steam

When the steam containing silica is injected to the turbine blades, it causes reduction of turbine efficiency by deposition on the turbine blades, nozzles etc and take part in corrosion.

Silica scale is mostly responsible for bulging and bursting of Wall Water Tubes and super heater tubes, because of its low thermal conductivity.

If silica is not removed from the boiler feed water, it will concentrate itself on the drum and is carried over in steam to form adherent deposits in the steam passage way distorting the original shape of turbine nozzles and blades. This alters steam velocities and the pressure drops reducing the capacity and efficiency of the turbine

Severe conditions can cause excessive rotor thrust while uneven deposition can unbalance the turbine rotor causing vibration problems. Turbine deposits can accumulate in a very short time, when steam purity is poor and can only removed by external service cleaning and blasting aluminium oxide on the surface

Experience has enabled the power industry to specify allowable concentrations of SiO2 in steam to avoid turbine damage. For a 180 bar operating pressure, in order to get a maximum of 5 ppb of SiO2 in the steam, the boiler water should not contain more than 100ppb of SiO2, if ideal boiler conditions are met.

Any minor deviations of silica concentrations in a power plant can have serious and expensive consequences in relation to performance, reliability, efficiency and safety. It is logical that silica concentrations should be monitored closely.

Method of Analysis

The chemistry employed for silica measurement is the industry standard Molybdenum Blue reaction.

The heteropoly blue method is used to measure molybdate-reactive silica. Molybdate 3 Reagent, an acidic molybdate solution, is added to the sample to react with any silica and phosphate present to form molybdosilicic and molybdophosphoric acids.

Citric Acid is then added, which masks any molybdophosphoic acid present and reacts with excess molybdate. This prevents molybdate from producing an interfering blue-colored compound. The surfactant, a wetting agent, minimizes air bubble formation on the sample-cell walls. Light absorbance through this solution is measured to determine a sample blank reference absorbance.

Color formed at this point is identical to the final color of a 0 µg/L silica sample. This provides a zero reference and compensates for any background turbidity and color inherent in the sample, changes in colorimeter lamp output or contamination of the sample-cell walls.

Amino Acid F Reagent is added to reduce molybdosilicic acid to a blue-colored solution. The amount of color formed is directly proportional to the silica concentration of the sample. Light absorbance through the solution is measured at 820 nm. This absorbance is compared to the sample-blank reference absorbance and the silica concentration is calculated.

Principle of Operation

One sample, appropriately conditioned (temperature, pressure), circulates at a relatively high linear velocity through its respective overflow sampling cup. A rate of 90 ml/hr for sample solution and 4.5 ml/hr of each of the 3 reagents is continuously being aspirated by means of a peristaltic pump and sequentially added into the analytical flow circuit. The high sample-to-reagent ratio minimizes errors arising as a result of inaccuracies of the reagent delivery rate by the pump. Figure. 1 highlights the flow diagram of typical Silica Analyzer.

The soluble silica of the sample now reacts with the molybdate and forms silicomolybdate. The rate of reaction of this complex formation is relatively slow. It is necessary, therefore, for the sample + molybdate to be mixed and kept in a reactor for 5 minutes in order to ensure completion of the reaction.

Although a direct photometric assay of the yellow complex is possible, the sensitivity of the measurement is insufficient for analysis of silica concentrations at the low ppb level. For this reason, the silicomolybdate complex is reduced with ferrous ions to the much more sensitive molydenumblue complex.

Prior to the addition of the reducing agent, the sample mix is reacted with oxalic acid to prevent interference by phosphates and at the same time, to intensify the color. The reducing agent (ferrous sulphate) is then added to the sample mix. The mix now enters the photometer flow cell where the optical density of the solution is measured by absorption of Infrared Light at a wavelength of 820 nm.

Refer Table 1.0 for explanation about parts and components highlighted in Figure. 1

Example of Actual Instrument ( Silica Analyzer)

Reagent Supply System

Reagents are supplied to the analysis module by pressurizing the reagent containers and using solenoid valves actuated by the control module to regulate flow volume and timing. Reagent containers are enclosed in a separate reagent compartment. A safety interlock on the compartment door requires reagent depressurization before opening. Reagent system pressure is supplied from an external source.

Analysis Module

The analysis module contains the solenoid valves containing sample and reagent flow and the colorimetric measurement system (Figure. 2). A sample-measurement cell (sample cell) is placed between a light source and a photo detector and filtered to measure light at 820 nm.

Analyzer Specifications

The following information is a typical specification of HACH Company’s Series 5000 Silica Analyzer:

• Range:
  0.00 to 5000 µg/L as SiO2
• Accuracy:
  0.00-5.00 µg/L : ±1.0 µg/L or ± 5% of reading, whichever is greater; 500-5000 µg/L; ±7% of reading
• Minimum Detection Limit: Less than 0.5 µg/L
• Precision: ±0.5 µg/L or ±1.0% of reading, whichever is greater
• Step Response Time: 8.8 minutes for 30 to 50 ºC, 15 minutes for 5 to 40 ºC (field adjustable)
• Ambient Operating Conditions :
  10 to 45 ºC, 5 to 95% non-condensing humidity. Suitable for general purpose, clean, indoor environments
• Analyzer Sample Requirements :
• Regulated to 5 ± 3 psig ( 34.5 ± 20.7 kPa). Flow rate from 100 to 300 mL/minute. Sample temperature between 5 and 50 ºC. A sample pressure control kit is provided
• Sample Inlet Fitting : ¼-inc OD stainless steel compression tubing fitting
• Recorder Outputs : Selectable for 0-0.01 V, 0-0.1 V, 0-1 V, or 4-20 mA. Output span programmable over any portion of 0-5000 µg/L range
• Serial I/O: RS232 and 20 mA current loop
• Alarms: Four programmable relays, two sample concentration alarms, analyzer system warning and analyzer system shutdown alarms each equipped with an SPDT relay, two with contacts rated for 1A resistive load at 30 VAC and 42 VDC and with two contacts rated for 5 A resistive load at 240 VAC
• Power Requirements: 115/230 VAC, 50/60 Hz, switch selectable; 52 VA, 32 W maximum
• Reagent Pressure Source : 20 to 60 psig regulated ( 137.9 to 413.7 kPa); nitrogen, instrument quality air or compressed air. Filter and regulator are supplied with analyzer
• Reagent Pressure Inlet Fitting : ¼-inch OD stainless steel compression tube fitting Sample Drain Fitting : ¾-inch NPT PVC Female
• Air Purge (Optional) : 5–scfh ( standard cubic feet per hour ) instrument quality air, ¼-inch OD stainless steel compression tube fitting
• Reagents : Reagents : 2.9 L Molubdate 3 , 2.9 L Citric Acid/Surfactant, 2.9 L Amino Acid F, 2.9 L Silica Standard Solution, SiO2, 500 µg/L ( 250 mL required for standardization)

Industrial Applications

The presence of silica in the steam and water circuits of power generation plant is associated with a number of problems both in the super heater and turbine sections. The solubility of silica in stream increases with pressure. The pressure of silica in the steam can lead to deposition in super heater tubes and on the turbine blades. Silica analyzer can give a very early warning if silica levels are going out of control and hence is very important analyzer.

Water Treatment Plant

Silica analyzers are used to measure the efficiency of the anion and mixed bed outlets in water treatment plant. The analyzers detect depletion of the beds and measure the final water quality to ensure it is suitable to enter the steam production cycle.

Boiler drum

Silica build-up is monitored insider the boiler drum. If the level of silica gets too high, then a ‘blow down’ is initiated to remove contaminated water from the boiler. Close control of silica levels will help minimize the frequency of boiler blow down, which can be expensive and inefficient if performed too often.

Boiler feed water

Measuring the levels of silica in boiler feed water will provide a final check on the quality and acceptability and will help ensure that the maximum permissible level of silica in the boiler is not exceeded.

Steam line

Silica monitoring within the steam line provides a good indication of the overall stream purity level provided by the boiler drum, ensuring the protection of the super heater and turbines

Benefits

Silica Analyzer- Troubleshooting

Most of the modern Silica Analyzers available in market is provided with self-diagnostic functions. Self diagnostic functions are used to detect certain types of system failures, actuate a system warning or alarm and display an error message describing the nature of the failure.

Problems with Consistency and Accuracy at Low Concentrations

Problems with consistent readings at lower concentrations may be caused by humidity in the environment. Humidity can condense on the sample cell wall in the light path if the sample temperature is below the dew point of the air next to the sample cell in the colorimeter.

Environment humidity and temperature can change throughout the day and from day to day. Humidity in a enclosed water plant may change when the building is closed up for cold weather or opened up for warm weather.

Actions to reduce potential humidity and temperature issues:

• Make sure the sample cell cover is tight.
• Seal any fittings that might leak fluids into the instrument
• Purge the instrument with dry instrument air or dry nitrogen to prevent excess humidity build up inside the instrument enclosure
• Place the instrument in an environmentally controlled ( temperature and humidity) building

Control of Silica in feed water

The concentration of silica (SiO2) in boiler feed water and boiler drum evaporation section should be controlled strictly to maintain minimum silica level in steam.

• Carryover of SiO2 in steam due to faulty operation should be avoided by maintaining accurate boiler drum level.
• Load should be increased gradually avoiding overloading, steam separator should be efficient, priming and foaming should be in low-pressure boiler
• Periodical and continuous blow down should be controlled strictly to maintain minimum level of TDS (Total Dissolved Solids ) in boiler water

Summary

Silica is a major culprit behind the build -up of hard and dense scale inside the boilers and turbines of power generation plants. At a time, when power companies are anxious to optimize their operations, in line with business and environmental pressures, they can ill afford to operate plants suffering from the impaired heat transfer that results from this type of fouling.

There’s an old saying in industry that ‘you can’t control what you don’t measure’ and silica in power generation is no exception. Silica deposits can impair the performance of equipment to such an extent that it is imperative to keep it under tight control.

The only way to effectively control silica build-up is through effective online silica monitoring analyzers, which provides an early warning of equipment problems before actual failure occurs, thereby ensuring the plant operates at best possible efficiency.