2. 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.

3. Show how [H+] = 1.0 X 10-7

4. Degree of ionisation is extremely small
Kw = Kc[H2O]= [H+][OH-] = = 1 x (at 25ºC)
Kw is the Ionic Product of water/dissociation product of water
Kw is temperature dependent ( not pressure or concentration dependent)
Increase temperature will increases the ionic product ( no effect on pH of water though)
Acidic solution [H+] greater [OH-]
Pure Water is a very very weak electrolyte
( only 1 in every 600 million water molecules ionise)

5. Kw is temperature dependent
T (°C)
Kw (mol2/litre2)
0.114 x 10-14 10
0.293 x 10-14 20
0.681 x 10-14 25
1.008 x 10-14 30
1.471 x 10-14 40
2.916 x 10-14 50
5.476 x 10-14
Kw of pure water increases as the temperature increases
The ionic product of water is the product of the hydrogen and hydroxide ion
concentration in 1litre of water at 25 °C
Kw = [H+][OH-] = 1 × at 25 °C

6. pH [H+ ] x [OH- ] = 1 x 10-14 = [1 x 10-7 ] x [1 x 10-7 ]
The ionic product of water is the product of the hydrogen and hydroxide ion
concentration in 1litre of water at 25 °C
Kw = [H+][OH-] = 1 × at 25 °C
[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+ ]

7. 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.

8. pH of Common Substances
Acidic Neutral Basic

9. The pH Scale Each pH unit is 10 times as large as the previous one9
Each pH unit is 10 times as large as the previous one
A change of 2 pH units means 100 times more basic or acidic
x10
x100
Limitations
Doesn’t cover very HIGH concentration (pH above 10-1) or very low pH values (pH below 10-14)
Must be aqueous
Affected by temperature ( standard temperature is 25°C)

10. concentration in 1litre of water at 25 °C
The ionic product of water is the product of the hydrogen and hydroxide ion
concentration in 1litre of water at 25 °C

11. 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

12. 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.

13. 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

14. 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.

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

16. The larger the hydrogen
pH scale
[H+] > 10-7M, pH < 7
ACIDIC
[H+] < 10-7M, pH > 7
BASIC
[H+] = 10-7M, pH = 7
NEUTRAL
The pH Scale
The larger the hydrogen
Ion concentration
The smaller the pH,
The stronger the acid

17. The pH Scale Each pH unit is 10 times as large as the previous one17
Each pH unit is 10 times as large as the previous one
A change of 2 pH units means 100 times more basic or acidic
x10
x100
Limitations
Doesn’t cover very HIGH concentration (pH above 10-1) or very low pH values (pH below 10-14)
Must be aqueous
Affected by temperature ( standard temperature is 25°C)

18. 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

19. pH & Indicators pH= 7 at 25° C
pH = -Log10 [H+]
Defined as the negative log to
the base 10 of the molar Hydrogen ion
concentration in an aqueous
solution

20. pH of bases: pOH
pOH= -log10 [OH-]
pH + pOH = 14
pH= 14 - pOH

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. 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

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

24. What is the pH of a 0.1 molar soltion of NaOH (careful)
What is the pH of 0.05 molar solution of Co(OH)2 ( assume its fully dissociated )

25. Solving for [H+]
A solution has a pH of What is the Molarity of hydrogen ions in the solution?
pH = - log [H+]
8.5 = - log [H+]

26. Strong and Weak Acids/Bases
Strong acids/bases – 100% dissociation into ions
HCl NaOH
HNO3 KOH
H2SO4
Weak acids/bases – partial dissociation, both ions and molecules
CH3COOH NH3
Need to know equilibrium constant

27. pH calculations for Weak Acids and Weak Bases
[H+]= √ka×Macid
[OH-]= √kb×Mbase
For Weak Acids
pH = -Log10
For Weak Bases
pOH = Log10
pH = pOH

28. 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 pH =

29. Calculate the pH of a 1 molar ethanoic acid solution that is only 1
Calculate the pH of a 1 molar ethanoic acid solution that is only 1.4% ionised

30. Acid base indicators
Substances that change colour according to pH of solution
Most are weak acids or bases so must only be added in small amounts. The colour of the dissociated molecule is different to the colour of the undissociated molecule
Some indicators dissociate to form weak bases
InH=In- + H+
InOH = In+ + OH-
Chemical equilibrium alters whether in presence of acid or base

31. 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)
Methyl orange
HIn (red, Acid)= H+ + In- (yellow, Base)
In acid: the equilibrium lies to the ______ giving it a ___ colour
In base: the equilibrium lies to the ______ giving it a ___ colour
: dynamic equilibrium: apply a stress by adding or removing H+ ions will shift the equilibrium
The equilibrium will shift depending on whether H+ ions or OH- ions exist. Therefore causing a colour change

32. Draw rough trend graph Name of Indicator Approx Range Acid Colour
Lower pH
Base Colour
Higher pH
Methyl Orange
red
yellow
Litmus
5-8
blue
Phenolphthalein
8.3-10
colourless
pink

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

34. Choice of Indicator for Titration
Indicator must have a complete colour change in the steepest 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. Question NB to practise
Page 261, 262
Question NB to practise