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Self Ionisation of Water

Self Ionisation of Water. Water undergoes Self Ionisation H 2 O (l) ⇄ H + (aq) + OH - (aq) or H 2 O (l) + H 2 O (l) ⇄ H 3 O + (aq) + OH - (aq) The concentration of H + ions and OH- ions is extremely small.

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Self Ionisation of Water

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  1. Self Ionisation of Water Water undergoes Self Ionisation H2O(l) ⇄ H+(aq) + OH-(aq) or H2O(l) + H2O(l) ⇄ H3O+(aq) + OH-(aq) The concentration of H+ ions and OH- ions is extremely small. Because the equilibrium lies very much on the left hand side.

  2. Glossary Ionisation Ionic Product pH Logarithm Kw Indicator pH scale Strong/weak acids Strong/Weak bases pH Curve End-Point Dissociation Constant

  3. Ionic Product of Water H2O(l) ⇄ H+(aq) + OH-(aq) Kc = In the above expression, the value of [H2O] may be taken as having a constant value because the degree of ionisation is so small. Kc = Kc [H2O] = [H+] [OH-] Both Kc and [H2O] are constant values so Kw = Kc [H2O] = [H+] [OH-] Kw = [H+] [OH-] is the ionic product of water

  4. Kw is temperature dependent Kw of pure water decreases as the temperature increases

  5. Acid–Base Concentrations in Solutions

  6. Acid–Base Concentrations in Solutions 10-1 OH- H+ 10-7 concentration (moles/L) H+ OH- OH- H+ 10-14 [H+]<[OH-] [H+]>[OH-] [H+]=[OH-] acidic solution neutral solution basic solution

  7. pH Scale Soren Sorensen (1868 - 1939) The pH scale was invented by the Danish chemist Soren Sorensen to measure the acidity of beer in a brewery. The pH scale measured the concentration of hydrogen ions in solution. The more hydrogen ions, the stronger the acid.

  8. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 The pH Scale 1 1 2 2 3 3 4 4 5 5 6 6 7 8 9 9 10 10 11 11 12 12 13 14 Strong Acid Weak Acid Neutral Weak Alkali Strong Alkali

  9. pH Scale The quantity of hydrogen ions in solution can affect the color of certain dyes found in nature. These dyes can be used as indicators to test for acids and alkalis. An indicator such as litmus (obtained from lichen) is red in acid. If base is slowly added, the litmus will turn blue when the acid has been neutralized, at about 6-7 on the pH scale. Other indicators will change color at different pH’s. A combination of indicators is used to make a universal indicator.

  10. Measuring pH • Universal Indicator Paper • Universal Indicator Solution • pH meter

  11. Measuring pH pH can be measured in several ways • Usually it is measured with a coloured acid-base indicator or a pH meter • Coloured indicators are a crude measure of pH, but are useful in certain applications • pH meters are more accurate, but they must be calibrated prior to use with a solution of known pH

  12. Limitations of pH Scale The pH scale ranges from 0 to 14 Values outside this range are possible but do not tend to be accurate because even strong acids and bases do not dissociate completely in highly concentrated solutions. pH is confined to dilute aqueous solutions

  13. pH At 250C Kw = 1 x 10-14 mol2/litre2 [H+ ] x [OH- ] = 1 x 10-14 mol2/litre2 This equilibrium constant is very important because it applies to all aqueous solutions - acids, bases, salts, and non-electrolytes - not just to pure water.

  14. pH For H2O(l)⇄H+(aq) + OH-(aq) → [H+ ] = [OH- ] [H+ ] x [OH- ] = 1 x 10-14 = [1 x 10-7 ] x [1 x 10-7 ] [H+ ] of water is at 250C is 1 x 10-7 mol/litre Replacing [H+ ] with pH to indicate acidity of solutions pH 7 replaces [H+ ] of 1 x 10-7 mol/litre where pH = - Log10[H+ ]

  15. pH is temperature dependent pH of pure water decreases as the temperature increases A word of warning! If the pH falls as temperature increases, does this mean that water becomes more acidic at higher temperatures? NO! Remember a solution is acidic if there is an excess of hydrogen ions over hydroxide ions. In the case of pure water, there are always the same number of hydrogen ions and hydroxide ions. This means that the water is always neutral - even if its pH change

  16. Students should be able to: • define pH • describe the use of the pH scale as a measure of the degree of acidity/alkalinity • discuss the limitations of the pH scale • explain self-ionisation of water • write an expression for Kw

  17. Acid – Base Concentrations and pH 10-1 pH = 11 pH = 3 OH- H+ pH = 7 10-7 concentration (moles/L) H+ OH- OH- H+ 10-14 [H3O+]<[OH-] [H3O+]>[OH-] [H3O+]=[OH-] acidic solution neutral solution basic solution

  18. pH describes both [H+ ] and [OH- ] 0 Acidic [H+ ] = 100[OH- ] =10-14 pH = 0 pOH = 14 • Neutral[H+ ] = 10-7 [OH- ] =10-7 pH = 7 pOH = 7 • Basic[H+ ] = 10-14[OH- ] = 100 pH = 14 pOH = 0

  19. pH of Common Substances Acidic Neutral Basic

  20. pH [H+] [OH-] pOH 14 1 x 10-14 1 x 10-0 0 13 1 x 10-13 1 x 10-1 1 12 1 x 10-12 1 x 10-2 2 11 1 x 10-11 1 x 10-3 3 10 1 x 10-10 1 x 10-4 4 9 1 x 10-9 1 x 10-5 5 8 1 x 10-8 1 x 10-6 6 6 1 x 10-6 1 x 10-8 8 5 1 x 10-5 1 x 10-9 9 4 1 x 10-4 1 x 10-10 10 3 1 x 10-3 1 x 10-11 11 2 1 x 10-2 1 x 10-12 12 1 1 x 10-1 1 x 10-13 13 0 1 x 100 1 x 10-14 14 NaOH, 0.1 M Household bleach Household ammonia Lime water Milk of magnesia Borax Baking soda Egg white, seawater Human blood, tears Milk Saliva Rain Black coffee Banana Tomatoes Wine Cola, vinegar Lemon juice Gastric juice More basic 7 1 x 10-7 1 x 10-7 7 More acidic

  21. You will need to memorize the following: Calculations and practice [H+] = 10–pH [OH–] = 10–pOH pH = –log10[H+] pOH = –log10[OH–] pH + pOH = 14

  22. pH Calculations pH [H+] pH = -log10[H+] [H+] = 10-pH pH + pOH = 14 [H+] [OH-] = 1 x10-14 pOH [OH-] pOH = -log10[OH-] [OH-] = 10-pOH

  23. Strong acids are so named because they react completely with water, leaving no undissociated molecules in solution. pH for Strong Acids • Strong acids dissociate completely in solution • Strong alkalis (bases) also dissociate completely in solution. • It is easy to calculate the pH of strong acids and strong bases; you only need to know the concentration.

  24. c) pH of solution where [H +] is 7.2x10-8M pH = – log10 [H+] = – log10 [7.2x10-8] = 7.14 (slightly basic) pH Exercises a) pH of 0.02M HCl pH = – log10 [H+] = – log10 [0.020] = 1.6989 = 1.70 b) pH of 0.0050M NaOH pOH = – log10 [OH–] = – log10 [0.0050] = 2.3 pH = 14 – pOH = 14 – 2.3 =11.7

  25. pH of dilute aqueous solutions of strong acids pH = ? pH = - log10 [H+] H1+(aq) + A1-(aq) HA(aq) monoprotic 0.3 M 0.3 M 0.3 M pH = - log10[0.3M] e.g. HCl, HNO3 pH = 0.48 pH = - log10[H+] 2 H1+(aq) + A2-(aq) H2A(aq) diprotic 0.3 M 0.6 M 0.3 M pH = - log10[0.6M] e.g. H2SO4 pH = 0.78

  26. 10x log 2nd antilog pH = - log [H+] Given: determine the [hydrogen ion] pH = 4.6 choose proper equation pH = - log10 [H+] substitute pH value in equation 4.6 = - log10 [H+] multiply both sides by -1 - 4.6 = log10[H+] take antilog of both sides - 4.6 = antilog [H+] [H+] = 2.51x10-5M You can check your answer by working backwards. pH = - log10[H+] pH = - log10[2.51x10-5M] pH = 4.6

  27. Most substances that are acidic in water are actually weak acids. Because weak acids dissociate only partially in aqueous solution, an equilibrium is formed between the acid and its ions. The ionization equilibrium is given by: HX(aq) H+(aq) + X-(aq) where X- is the conjugate base.

  28. pH calculations for Weak Acids and Weak Bases For Weak Acids pH = -Log10 For Weak Bases pOH = Log10 pH = 14 - pOH

  29. Calculating pH - weak acids A weak acid is one which only partially dissociates in aqueous solution A weak acid, HA, dissociates as followsHA(aq) H+(aq) + A¯(aq)(1) Applying the Equilibrium Law Ka = [H+(aq)] [A¯(aq)] mol dm-3(2) [HA(aq)] The ions are formed in equal amounts, so [H+(aq)] = [A¯(aq)] therefore Ka = [H+(aq)]2(3) [HA(aq)] Rearranging (3) gives[H+(aq)]2 = [HA(aq)]Ka therefore [H+(aq)] = [HA(aq)]Ka

  30. pH of solutions of weak concentrations Weak Acid pH of a 1M solution of ethanoic acid with a Ka value of 1.8 x 10-5 pH = -Log10 pH = -Log10 pH = 2.3723

  31. pH of solutions of weak concentrations Weak Base pH of a 0.2M solution of ammonia with a Kb value of 1.8 x 10-5 pOH = -log10 pOH = -log10 pOH = 2.7319 pH = 14 – 2.7319 pH = 11.2681

  32. Theory of Acid Base Indicators Acid-base titration indicators are quite often weak acids. For the indicator HIn The equilibrium can be simply expressed as HIn(aq, colour 1) H+(aq) + In-(aq, colour 2) The un-ionised form (HIn) is a different colour to the anionic form (In¯).

  33. Theory of Acid Base Indicators • Applying Le Chatelier's equilibrium principle: • Addition of acid • favours the formation of more HIn (colour 1) • HIn(aq) H+(aq) + In-(aq) • because an increase on the right of [H+] • causes a shift to left • increasing [HIn] (colour 1) • to minimise 'enforced' rise in [H+].

  34. Theory of Acid Base Indicators • Applying Le Chatelier's equilibrium principle: • Addition of base • favours the formation of more In- (colour 2) • HIn(aq) H+(aq) + In-(aq) • The increase in [OH-] causes a shift to right because the reaction H+(aq) + OH-(aq) ==> H2O(l) • Reducing the [H+] on the right • so more HIn ionises to replace the [H+] • and so increasing In- (colour 2) • to minimise 'enforced' rise in [OH-]

  35. Theory of Acid Base Indicators • Summary • In acidic solution • HIn(aq)H+(aq) + In¯(aq) • In alkaline solution

  36. Theory of Acid Base Indicators Acid-base titration indicators are also often weak bases. For the indicator MOH The equilibrium can be simply expressed as MOH(aq, colour 1) OH-(aq) + M+(aq, colour 2)

  37. Theory of Acid Base Indicators • Applying Le Chatelier's equilibrium principle: • Addition of base • favours the formation of more MOH (colour 1) • MOH(aq) M+(aq) + OH-(aq) • because an increase on the right of [OH-] • causes a shift to left • increasing [MOH] (colour 1) • to minimise 'enforced' rise in [OH-].

  38. Theory of Acid Base Indicators • Applying Le Chatelier's equilibrium principle: • Addition of acid • favours the formation of more M+ (colour 2) • MOH(aq) M+(aq) + OH-(aq) • The increase in [H+] causes a shift to right because the reaction H+(aq) + OH-(aq) ==> H2O(l) • Reducing the [OH-] on the right • so more MOH ionises to replace the [OH-] • and so increasing M+ (colour 2) • to minimise 'enforced' rise in [H+]

  39. 25 cm3 of 0.1 mol dm-3 acid is titrated with 0.1 mol dm-3 alkaline solution. Acid Base Titration Curves Strong Acid – Weak Base Strong Acid – Strong Base Weak Acid – Weak Base Weak Acid – Strong Base

  40. Choice of Indicator for Titration • Indicator must have a complete colour change in the vertical part of the pH titration curve • Indicator must have a distinct colour change • Indicator must have a sharp colour change

  41. Indicators for Strong Acid Strong Base Titration Both phenolphthalein and methyl orange have a complete colour change in the vertical section of the pH titration curve

  42. Indicators for Strong Acid Weak Base Titration Methyl Orange is used as indicator for this titration Only methyl orange has a complete colour change in the vertical section of the pH titration curve Phenolphthalein has not a complete colour change in the vertical section on the pH titration curve.

  43. Indicators for Weak Acid Strong Base Titration Phenolphthalein is used as indicator for this titration Only phenolphthalein has a complete colour change in the vertical section of the pH titration curve Methyl has not a complete colour change in the vertical section on the pH titration curve.

  44. Indicators for Weak Acid Weak Base Titration No indicator suitable for this titration because no vertical section Neither phenolphthalein nor methyl orange have completely change colour in the vertical section on the pH titration curve

  45. Colour Changes and pH ranges

  46. Methyl Orange

  47. Phenolphthalein

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