Titrimetric Analysis

1. Titrimetric Analysis • 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 • Neutralisation Reactions • Complex Formation Reactions • Redox Reactions • Precipitation Reactions

3. Basics • Equivalence and end points • Standards

4. Basics • Equivalence and end points • 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.

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

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

7. Basics • An indicator constant is defined as: KIn = [InB][X]n / [InA] [X]n = KIn ([InB] / [InA]) npX = pKIn + log10([InB] / [InA]) pH = pKa + log10 ([InB] / [InA])

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

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

10. Basics

11. Basics [X]x[Y]y... [A]a[B]b... RT nF E = E0 - ln • The Nernst Equation aA + bB + …+ ne- = xX + yY + ...

12. Basics [X]x[Y]y... [A]a[B]b... E = E0 - (0.059 V/n) log10 • RT/F ln 10 = 0.059158 V thus: • And E = E0 at unity concentrations

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

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

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

16. Basics

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

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

19. Basics • Acid-base reactions. • Na2CO3, Na2B4O7, KH(C8H4O4), HCl (cbpt.) • Complex formation reactions. • AgNO3, NaCl • Precipitation reactions. • AgNO3, KCl • Redox reactions. • K2Cr2O7,Na2C2O4, I2

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

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

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

23. Neutralisation Titrations • For any actual titration the correct end point will be characterised by a definite value of the hydrogen ion concentration. • 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 • A large number of substances called neutra-lisation indicators change colour according to the hydrogen ion concentration of the solution. • The end point can also be determined electrochemically by either potentiometric or conductimetric methods.

25. Theory of Indicator Behaviour • 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. • The behaviour of an acid type indicator is described by the equilibrium;

26. Theory of Indicator Behaviour • HIn + H2O In- + H3O+ • In + H2O InH+ + OH- • The behaviour of an base type indicator is described by the equilibrium;

27. Theory of Indicator Behaviour [H3O+][In-] [HIn] = Ka [HIn-] [In-] [H3O+] = Ka • The equilibrium constant takes the form: • Rearranging:

28. Theory of Indicator Behaviour • pH (acid colour) = -log(Ka . 10) = pKa +1 • pH (base colour) = -log(Ka / 10) = pKa -1 • Therefore; indicator range = pKa ± 1

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

30. Theory of Indicator Behaviour

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

32. Neutralisation Titrations H3O+ + OH- 2H2O • Strong acids and bases. • When both reagent and analyte are strong electrolytes, the neutralisation reaction can be described by the equation:

33. Neutralisation Titrations • The H3O+ concentration in aqueous solution comprises of two components. • The reaction of the solute with water. • The dissociation of water. • [H3O+] = CHCl + [OH-] = CHCl • [OH-] = CNaOH + [H3O+] = CNaOH

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

35. Neutralisation Titrations

36. Neutralisation Titrations • Examples: • HCl, HNO3 • NaOH, KOH, Na2CO3 • Standards:anhydrous Na2CO3 and constant boiling HCl.

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

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

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

40. Neutralisation Titrations

41. Neutralisation Titrations • Polyfunctional acids and bases • Typified by more than one dissociation reaction.

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

43. Neutralisation Titrations

44. Neutralisation Titrations • Sulphuric Acid • Unusual because one proton behaves as a strong acid and the other as a weak acid (K2 = 1.20 x 10-2).

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

46. Complexometric Titrations • Titrations between cations and complex forming reagents. • 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 • Most metal ions react with electron-pair donors to form coordination compounds or complex ions. • The donor species, or LIGAND, must have at least one pair of unshared electrons available.

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

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

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