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. H2O(l) ⇄ H+(aq) + OH-(aq)
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 [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

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

4. pH Scale
Soren Sorensen
( )
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.
The pH scale was invented by the Danish chemist Soren Sorensen for a brewery to measure the acidity of beer. 4

5. The pH Scale 7 8 9 10 11 12 13 3 4 5 6 2 14 1 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

6. 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. 6

7. Measuring pH Universal Indicator Paper Universal Indicator Solution
pH meter

8. Measuring pH pH can be measured in several ways
30/09/99
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 8

9. 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

10. pH Kw = 1 x 10-14 mol2/litre2 [H+ ] x [OH- ] = 1 x 10-14 mol2/litre2
At 250C
Kw = 1 x mol2/litre2
[H+ ] x [OH- ] = 1 x 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.

11. [H+ ] of water is at 250C is 1 x 10-7 mol/litrepH
For H2O(l) ⇄ H+(aq) + OH-(aq)
→ [H+ ] = [OH- ]
[H+ ] x [OH- ] = 1 x = [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 mol/litre
where pH = - Log10 [H+ ]

12. pH is temperature dependent
T (°C) pH
7.12 10
7.06 20
7.02 25 7 30
6.99 40
6.97
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

13. Acid – Base Concentrations and pH
10-1
pH = 11
pH = 3 H+
OH-
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 13

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

16. pH of Common Substances
Acidic Neutral Basic 17

17. pH [H+] [OH-] pOH
14 1 x x
13 1 x x
12 1 x x
11 1 x x
10 1 x x
9 1 x x
8 1 x x
6 1 x x
5 1 x x
4 1 x x
3 1 x x
2 1 x x
1 1 x x
0 1 x x
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 x
More acidic 18

18. Calculations and practice
30/09/99
You will need to use the following:
pH = – log10[H+]
pOH = – log10[OH–]
pH + pOH = 14 19

19. pH Calculations pH pOH [H+] [OH-] pH = -log10[H+] [H+] = 10-pH
[H+] [OH-] = 1 x10-14
pOH
[OH-]
pOH = -log10[OH-]
[OH-] = 10-pOH 20

20. pH for Strong Acids Strong acids dissociate completely in solution
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 bases also dissociate completely in solution

21. pH Exercises
30/09/99
c) pH of solution where [H +] is 7.2x10-8M
pH = – log10 [H+]
= – log10 [7.2x10-8] =
(slightly basic)
a) pH of 0.02M HCl pH = – log10 [H+] = – log10 [0.020] = = 1.70 b) pH of M NaOH pOH = – log10 [OH–] = – log10 [0.0050] = 2.3 pH = 14 – pOH = 14 – 2.3 =11.7 22

22. pH = - log [H+] Given: pH = 4.6 determine the [hydrogen ion]
choose proper equation
4.6 = - log10 [H+]
substitute pH value in equation
= log10[H+]
multiply both sides by -1
2nd
log
= antilog [H+]
take antilog of both sides
[H+] = 2.51x10-5 M
10x
antilog
You can check your answer by working backwards.
pH = - log10[H+]
pH = - log10[2.51x10-5 M]
pH = 4.6 24

23. 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.

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

25. 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 =

26. 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 = pH = 14 – pH =

27. 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)

28. 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+].

29. 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-]

30. 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)

31. 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-].

32. 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+]

33. Acid Base Titration Curves
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 35

34. 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

35. 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

36. 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.

37. 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.

38. 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

39. indicator
pH range
litmus
5 - 8
methyl orange
phenolphthalein

40. Colour Changes and pH ranges

41. Methyl Orange

42. Phenolphthalein

43. Universal indicator components
Low pH color
Transition pH range
High pH color
Thymol blue (first transition)
red
1.2–2.8
orange
Methyl Orange
4.4–6.2
yellow
Bromothymol blue
6.0–7.6
blue
Thymol blue (second transition)
8.0–9.6
Phenolphthalein
colourless
8.3–10.0
purple