1. Titrimetric Analysis l Quantitative chemical analysis carried out by determining the volume of a solution of accurately known concentration which is required to react quantitatively with a measured volume of the substance to be determined.

2. Classification l Neutralisation Reactions l Complex Formation Reactions l Redox Reactions l Precipitation Reactions

3. Basics l Equivalence and end points l Standards

4. l Equivalence and end points v Precise and accurate titrations require the reproducible determination of the end point which either corresponds to the stoichiometric point of the reaction or bears a fixed and measurable relation to it. Basics

5. l Equivalence and end points v Monitor a property of the titrand which is removed at the end point. v Monitor a property which is readily observed when excess titrant has been added. v Two main methods n Coloured indicators n Electrochemical techniques. Basics

6. l Colour Change Indicators v Common to a wide variety of titrations. v In general terms a visual indicator is a compound that changes from one colour to another as its chemical form changes. v In A = In B + nX where X may be H +, M n+ or e -, and the colour is sensitive to the presence of H +, M n+, oxidants or reductants. Basics

7. l An indicator constant is defined as: K In = [In B ][X] n / [In A ] [X] n = K In ([In B ] / [In A ]) npX = pK In + log 10 ([In B ] / [In A ]) pH = pK a + log 10 ([In B ] / [In A ]) Basics

8. l Potentiometric Measurements l Measuring the change in potential during the titration. v Acid-base titrations. v Precipitation titrations. v Redox titrations. Basics

9. l Monitor the change of E cell during the course of a titration where the indicator electrode responds to one of the reactants or the products. l A plot of E cell against the volume of titrant is obtained. l Precision of better than 0.2%. Basics

10. Basics

11. l The Nernst Equation aA + bB + …+ ne - = xX + yY +... E = E 0 - ln RT nF [X] x [Y] y... [A] a [B] b...Basics

12. l RT/F ln 10 = 0.059158 V thus: E = E 0 - (0.059 V/n) log 10 [X] x [Y] y... [A] a [B] b... l And E = E 0 at unity concentrations Basics

13. l Conductimetric Indication v The electrical conductance of a solution is a measure of its current carrying capacity and is determined by its total ionic strength. v It is a non-specific property.  Conductance is defined as the reciprocal of resistance (Siemans,  -1 ). Basics

14. l A conductance cell consists of two platinum electrodes of large surface area. l 5-10 V at 50 -10,000 Hz is applied. l Control of temperature is essential. Basics

15. l Acid-base titrations especially at trace levels. l Relative precision better than 1% at all levels. l Rate of change of conductance as a function of added titrant used to determine the equivalence point. l High concentrations of other electrolytes can interfere. Basics

16. Basics

17. l Standards v Certain chemicals which are used in defined concentrations as reference materials. v Primary standards. v Secondary standards. Basics

18. l Primary Standards v Available in pure form, stable and easily dried to a constant known composition. v Stable in air. v High molecular weight. v Readily soluble. v Undergoes stoichiometric and rapid reactions. Basics

19. v Acid-base reactions. n Na 2 CO 3, Na 2 B 4 O 7, KH(C 8 H 4 O 4 ), HCl (cbpt.) v Complex formation reactions. n AgNO 3, NaCl v Precipitation reactions. n AgNO 3, KCl v Redox reactions. n K 2 Cr 2 O 7, Na 2 C 2 O 4, I 2 Basics

20. l Secondary Standards v A substance that can be used for standardisations, and whose concentration of active substance has been determined by comparison to a primary standard. Basics

21. Classification l Neutralisation Reactions l Complex Formation Reactions l Redox Reactions l Precipitation Reactions

22. Neutralisation Titrations l The neutralisation reactions between acids and bases used in chemical analysis. l These reactions involve the combination of hydrogen and hydroxide ions to form water.

23. Neutralisation Titrations l For any actual titration the correct end point will be characterised by a definite value of the hydrogen ion concentration. l This value will depend upon the nature of the acid and the base, the concentration of the solution and the nature of the indicator.

24. Neutralisation Titrations l A large number of substances called neutra- lisation indicators change colour according to the hydrogen ion concentration of the solution. l The end point can also be determined electrochemically by either potentiometric or conductimetric methods.

25. l An acid/base indicator is a weak organic acid or a weak organic base whose undissociated form differs in colour from its conjugate base or conjugate acid form. l The behaviour of an acid type indicator is described by the equilibrium; Theory of Indicator Behaviour

26. l The behaviour of an base type indicator is described by the equilibrium; l In + H 2 O InH + + OH - l HIn + H 2 O In - + H 3 O +

27. l The equilibrium constant takes the form: [H 3 O + ][In - ] [HIn] = K a [H 3 O + ]= K a [HIn - ] [In - ] Theory of Indicator Behaviour l Rearranging:

28. l pH (acid colour) = -log(K a. 10) = pK a +1 l pH (base colour) = -log(K a / 10) = pK a -1 l Therefore; indicator range = pK a ± 1 Theory of Indicator Behaviour

29. l The human eye is not very sensitive to colour change in a solution containing In - and HIn. l Especially when the ratio [In - ] / [HIn] is greater than 10 or less than 0.1. l Hence the colour change is only rapid within the limited concentration ratio of 10 to 0.1. Theory of Indicator Behaviour

31. Neutralisation Titrations l Strong acids and bases l Weak acids l Weak bases l Polyfunctional acids l Applications

32. Neutralisation Titrations l Strong acids and bases. l When both reagent and analyte are strong electrolytes, the neutralisation reaction can be described by the equation: H 3 O + + OH - 2H 2 O

33. Neutralisation Titrations l The H 3 O + concentration in aqueous solution comprises of two components. v The reaction of the solute with water. v The dissociation of water. l [H 3 O + ] = C HCl + [OH - ] = C HCl l [OH - ] = C NaOH + [H 3 O + ] = C NaOH

34. Neutralisation Titrations l Using these assumptions you can calculate the pH of a titration solution directly from stoichiometric calculations and therefore simulate the titration curves. l This is useful in determining the correct indicator for a new titration.

35. Neutralisation Titrations

36. l Examples: v HCl, HNO 3 v NaOH, KOH, Na 2 CO 3 v Standards:anhydrous Na 2 CO 3 and constant boiling HCl.

37. Neutralisation Titrations l Weak acids and bases l Examples v Ethanoic acid v Sodium cyanide l Four types of calculation are required to derive a titration curve for a weak acid or base.

38. Neutralisation Titrations v Solution contains only weak acid. pH is calculated from the concentration and the dissociation constant. v After additions of the titrant the solution behaves as a buffer. The pH of each buffer can be calculated from there analytical concentrations. v At the equivilence point only salt is present and the pH is calculated from the concentration of this product. v Beyond the equivilence point the pH is governed largely by the concentration of the excess titrant.

39. Neutralisation Titrations l Effect of Concentration l Effect of reaction completeness l Indicator choice; Feasibility of titration

40. Neutralisation Titrations

41. l Polyfunctional acids and bases l Typified by more than one dissociation reaction.

42. Neutralisation Titrations l Phosphoric acid l Yield multiple end points in a titration.

43. Neutralisation Titrations

44. l Sulphuric Acid l Unusual because one proton behaves as a strong acid and the other as a weak acid (K 2 = 1.20 x 10 -2 ).

45. Neutralisation Titrations l Applications: v Determination of the concentration of analytes which are either acid or bases. v Determination of analytes which can be converted to acids or bases.

46. Complexometric Titrations l Titrations between cations and complex forming reagents. l The most useful of these complexing agents are organic compounds with several electron donor groups that can form multiple covalent bonds with metal ions.

47. Complexometric Titrations l Most metal ions react with electron-pair donors to form coordination compounds or complex ions. l The donor species, or LIGAND, must have at least one pair of unshared electrons available.

48. Complexometric Titrations l Inorganic Ligands v Water v Ammonia v Halides l Organic Ligands v Cyanide v Acetate

49. Complexometric Titrations l The number of bonds a cation forms with an electron donor is called the COORDINATION NUMBER. l Typical values are 2, 4 and 6. l The species formed as a result of coordination can be electrically positive, neutral or negative.

50. Complexometric Titrations l Complexometric methods have been around for more than a century. l Rapid expansion in the 1940’s based on a class of coordination compounds called CHELATES.

51. Complexometric Titrations l A chelate is produced when a metal ion coordinates to two or more donor groups within a single ligand. l For example the copper complex of glycine.

52. Complexometric Titrations Cu 2+ + 2 HC NH 2 H C O OH O N H 2 C C H 2 O Cu O N H 2 C C H 2 O + 2H +

53. Complexometric Titrations l A ligand with a single donor group is called unidentate. l Glycine is bidentate. l Tri, tetra, penta and hexadentate chelating agents are also known.

54. Complexometric Titrations l Multidentate ligands have two advantages over unidentate ligands. l They react more completely with cations to provide a sharper endpoint. l The reaction is a single step process.

55. Complexometric Titrations l Tertiary amines that also contain carboxylic acid groups form remarkably stable chelates with many metal ions. l Ethylenediaminetetraacetic Acid EDTA CH 2 CH 2 N N HOOCCH 2 CH 2 COOH CH 2 COOH HOOCCH 2

56. Complexometric Titrations l EDTA can complex a large number of metal ions. l Approximately 40 cations can be determined by direct titration. l EDTA is usually used as the disodium salt, Na 2 H 2 EDTA H 2 EDTA 2- + M 2+  [M(EDTA)] 2- + 2H +

57. Complexometric Titrations l Because EDTA complexes most cations, the reagent might appear at first glance to be totally lacking in selectivity. l However, great control can be acheived by pH regulation and the selection of suitable indicators.

58. Complexometric Titrations l Indicators are generally complexing agents which undergo a colour change when bonded to a metal ion. H 2 EDTA 2- + [M(Ind)]  [M(EDTA)] 2- + Ind 2- + 2H +

59. Complexometric Titrations l Typical indicators are: v Murexide v Solochrome black v Calmagite v Bromopyrogallol red v Xylenol orange

61. Complexometric Titrations l Typical applications: v Determination of cations v Hardness of water

62. Redox Titrations l Basics l Potassium Permanganate l Potassium Dichromate l Cerium IV l Iodine

63. Redox Titrations l Basics v Electrode Potentials v Indicators

64. Redox Titrations l Electrode Potentials v Derived from Nernst equation. v Calculations of cell potentials leads to theoretical titration curves. v E OX = E RED = E system

65. Redox Titrations

66. l Indicators v Potentiometric v Coloured indicators v E In = E OX = E RED = E system v Specific: Starch v Oxidation / Reduction Indicators

67. Redox Titrations

68. l Potassium Permanganate MnO 4 - + 8H + + 5e -  Mn 2+ + 4H 2 O l Standardisation v Sodium oxalate or arsenic (III) oxide l Many Analyses

69. Redox Titrations l Hydrogen Peroxide:  2MnO 4 - + 5H 2 O 2 + 6H +  2Mn 2+ + 5O 2 + 8H 2 O l Nitrites:  2MnO 4 - + 5NO 2 - + 6H +  2Mn 2+ + 5NO 3 - + 3H 2 O

70. Redox Titrations l Persulphates: v Add an excess of iron (II)  S 2 O 8 2- + 2Fe 2+ + 2H +  2Fe 3+ + 2HSO 4 - v The excess iron (II) is determined by back titration against standardised permangenate.  MnO 4 - + 8H + + 5Fe 2+  Mn 2+ + 5Fe 2+ + 4H 2 O

71. Redox Titrations l Potassium Dichromate CrO 7 2- + 14H + + 6e -  2Cr 3+ + 7H 2 O l Standardisation v Against metallic iron v 1 mole K 2 CrO 7 = 6 moles Fe

72. Redox Titrations l Iron (II): CrO 7 2- + 14H + + 6Fe 2+  2Cr 3+ + 6Fe 3+ 7H 2 O v Indicators include N-phenylanthranilic acid and sodium diphenylamine sulphonate. l Chlorates: v Reduced with an excess of iron (II) v The excess iron (II) is determined by back titration against standardised dichromate.

73. Redox Titrations l Iodine l Iodometric Titrations I 2 + 2e -  2I - l Standardisation v Standardised sodium thiosulphate or arsenic (III) oxide l Many Analyses

74. Redox Titrations l Hydrogen Peroxide:  H 2 O 2 + 2I - + 2H +  I 2 + 2H 2 O l Thiosulphates:  2S 2 O 3 2- + I 2  S 4 O 6 2- + 2I - l Hydroxyl Groups:  2OH - + I 2  IO - + H 2 O + 2I -

75. Redox Titrations l Others: v Copper v Dissolved oxygen v Chlorine v Arsenic (V) v Sulphides v etc..........

76. Redox Titrations l Cerium (IV) Sulphate l Very strong oxidising agent (1.43V) Ce 4+ + e -  Ce 3+ l Standardisation v Sodium oxalate or arsenic (III) oxide l Many Analyses

77. Precipitation Titrations l Titrations between analytes and reagents resulting in the formation of a precipitate. l The most useful of these precipitating reagents is silver nitrate. l Titrimetric methods based upon the use of silver nitrate are sometimes called Argentometric titrations.

78. Precipitation Titrations l Used for the determination of many anions including: v halides v divalent anions v mercaptans v certain fatty acids

79. Precipitation Titrations l Precipitation titrations are based on the SOLUBILITY PRODUCT of the salt, K SP. l The smaller K SP, the less soluble the silver salt and the easier it is to determine the endpoint

80. Precipitation Titrations l Endpoint determination is by coloured indicators (usually back titrations) or turbidity methods. l The most accurate is the VOLHARD METHOD.

81. Precipitation Titrations l VOLHARD METHOD l A back titration of thiocyanate ions against the excess silver ions using an iron (II) salt as the indicator.

82. Precipitation Titrations Blood Red Ag + + SCN - AgSCN Fe 3+ + SCN - FeSCN 2+