Self Ionisation of Water

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
T (C) Kw (mol2/litre2)
0 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 decreases as the temperature
increases
5
AcidBase Concentrations in Solutions
6
AcidBase Concentrations in Solutions
10-1
OH-
H
10-7
concentration (moles/L)
OH-
H
OH-
H
10-14
H lt OH-
H OH-
H gt 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
The pH Scale
7
8
9
10
11
12
13
3
4
5
6
2
14
1
9
10
11
12
3
4
5
6
2
1
Neutral
Weak Alkali
Strong Alkali
Weak Acid
Strong Acid
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 pHs. 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
T (C) pH
0 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
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)
OH-
H
OH-
H
10-14
H3O lt OH-
H3O OH-
H3O gt 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
14
pH -log10 H
20

21
pH of Common Substances
Acidic
Neutral
Basic
22
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 Bakin
g soda Egg white, seawater Human blood,
tears Milk Saliva Rain Black coffee Banana Tomato
es Wine Cola, vinegar Lemon juice Gastric juice
More basic
7 1 x 10-7 1 x 10-7 7
More acidic
23
Calculations and practice

• You will need to memorize the following

H 10pH
OH 10pOH
pH log10H
pOH log10OH
pH pOH 14
24
pH Calculations
pH
H
pH -log10H
H 10-pH
pH pOH 14
H OH- 1 x10-14
pOH
OH-
pOH -log10OH-
OH- 10-pOH
25
pH for Strong Acids
Strong acids are so named because they react
completely with water, leaving no undissociated
molecules in solution.

• 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.

26
pH Exercises
c) pH of solution where H is 7.2x10-8M pH
log10 H log10 7.2x10-8
7.14 (slightly basic)

• 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

27
OH-
1.0 x 10-7 OH-
10-pOH
1.0 x 10-7 H
-Log10OH-
H
pOH
10-pH
14 - pOH
-Log10H
pH
pH
28
pH and pOH

• pH - log10H3O H3O 10-pH
• pOH - log10OH- OH- 10-pOH
• pKw pH pOH 14.00
• neutral solution H3O OH- 10 7 M pH
  7.0
• acidic solution H3O gt 10-7 M pH lt
  7.0
• basic solution H3O lt 10-7 M pH gt 7.0

29
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 - log100.3M
e.g. HCl, HNO3
pH 0.48
pH - log10H
2 H1(aq) A2-(aq)
H2A(aq)
diprotic
0.3 M
0.6 M
0.3 M
pH - log100.6M
e.g. H2SO4
pH 0.78
30
A sample of orange juice has a hydrogen-ion
concentration of 2.9 x 10-4M. What is the pH?
pH -log10 H pH -log10 (2.9x10-4 ) pH
3.54
31
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 log10H
take antilog of both sides
- 4.6 antilog H
H 2.51x10-5 M
You can check your answer by working backwards.
pH - log10H
pH - log102.51x10-5 M
pH 4.6
32
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.
33
pH calculations for Weak Acids and Weak Bases
For Weak Acids pH -Log10
For Weak Bases
pOH Log10
pH 14 - pOH
34
Calculating pH - weak acids
A weak acid is one which only partially
dissociates in aqueous solution
A weak acid, HA, dissociates as follows HA(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

35
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

36
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

37
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).

38
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.

39
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) gt 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-
40
Theory of Acid Base Indicators

• Summary
• In acidic
  solution
• HIn(aq) H(aq) In(aq)
• In alkaline solution

41
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)

42
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-.

43
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) gt 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
44
Acid Base Titration Curves
25 cm3 of 0.1 mol dm-3 acid is titrated with 0.1
mol dm-3 alkaline solution.
Strong Acid Weak Base
Strong Acid Strong Base
Weak Acid Strong Base
45
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

46
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
47
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.
48
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.
49
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
50
indicator pH range
litmus 5 - 8
methyl orange 3.1 - 4.4
phenolphthalein 8.3 - 10.0
51
Colour Changes and pH ranges
52
Methyl Orange
53
Phenolphthalein
54
Universal indicator components Universal indicator components Universal indicator components Universal indicator components
Indicator Low pH color Transition pH range High pH color
Thymol blue (first transition) red 1.22.8 orange
Methyl Orange red 4.46.2 yellow
Bromothymol blue yellow 6.07.6 blue
Thymol blue (second transition) yellow 8.09.6 blue
Phenolphthalein colourless 8.310.0 purple
55
Students should be able to

• calculate the pH of dilute aqueous solutions of
  strong acids and bases
•  
• distinguish between the terms weak, strong,
  concentrated and dilute in relation to acids and
  bases
•  
• calculate the pH of weak acids and bases
  (approximate method of calculation to be used
  assuming that ionisation does not alter the total
  concentration of the non-ionised form)
•  
• define acid-base indicator
•  
• explain the theory of acid-base indicators
•  
• justify the selection of an indicator for acid
  base titrations