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.
Show how [H+] = 1.0 X 10-7
• Degree of ionisation is extremely small
• Kw = Kc[H2O]= [H+][OH-] = = 1 x 10-14 (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)
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 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 × 10-14 at 25 °C
pH
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 × 10-14 at 25 °C
[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+ ]
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.
pH of Common Substances
Acidic
Neutral
Basic
9
The pH Scale
• 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
1. Doesn’t cover very HIGH concentration (pH above 10-1) or
very low pH values (pH below 10-14)
2. Must be aqueous
3. Affected by temperature ( standard temperature is 25°C)
The ionic product of water is the product of the hydrogen and hydroxide ion
concentration in 1litre of water at 25 °C
Acid–Base Concentrations in Solutions
concentration (moles/L)
10-1
OH-
H+
10-7
H+
OH-
OH-
H+
10-14
[H+] > [OH-]
[H+] = [OH-]
acidic
solution
neutral
solution
[H+] < [OH-]
basic
solution
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.
The pH Scale
1
2
Strong
Acid
3
4
Weak
Acid
5
6
7
Neutral
8
9
10
11
Weak
Alkali
12
13
14
Strong
Alkali
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.
Measuring pH
• Universal Indicator Paper
• Universal Indicator Solution
• pH meter
pH scale
[H+] > 10-7M, pH < 7
ACIDIC
[H+] < 10-7M, pH > 7
BASIC
[H+] = 10-7M, pH = 7
NEUTRAL
The larger the hydrogen
Ion concentration
The smaller the pH,
The stronger the acid
The pH Scale
The pH Scale
17
• 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
1. Doesn’t cover very HIGH concentration (pH above 10-1) or
very low pH values (pH below 10-14)
2. Must be aqueous
3. Affected by temperature ( standard temperature is 25°C)
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
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
pH of bases: pOH
pOH= -log10 [OH-]
pH + pOH = 14
pH= 14 - pOH
pH Exercises
c) pH of solution where [H +]
a)pH of 0.02M HCl
is 7.2x10-8M
+
pH = – log10 [H ]
pH
= – log10 [H+]
= – log10 [0.020]
= – log10 [7.2x10-8]
= 1.6989
= 7.14
= 1.70
(slightly basic)
b)pH of 0.0050M NaOH
pOH = – log10 [OH–]
= – log10 [0.0050]
= 2.3
pH = 14 – pOH
= 14 – 2.3
=11.7
pH Calculations
pH
pH = -log10[H+]
[H+]
[H+] = 10-pH
[H+] [OH-] = 1 x10-14
pH + pOH = 14
pOH
pOH = -log10[OH-]
[OH-]
[OH-] = 10-pOH
pH of dilute aqueous solutions of acids
monoprotic
e.g. HCl, HNO3
HA(aq)
0.3 M
H1+(aq) + A1-(aq)
0.3 M
0.3 M
pH = ?
pH = - log10 [H+]
pH = - log10[0.3M]
pH =
diprotic
e.g. H2SO4
H2A(aq)
0.3 M
2 H1+(aq) + A2-(aq)
0.6 M
0.3 M
0.52
pH = - log10[H+]
pH = - log10[0.6M]
pH =
0.22
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 )
Solving for [H+]
• A solution has a pH of 8.5. What is the
Molarity of hydrogen ions in the
solution?
pH = - log [H+]
8.5 = - log [H+]
Strong and Weak Acids/Bases
Strong acids/bases – 100% dissociation into ions
HCl
HNO3
H2SO4
NaOH
KOH
Weak acids/bases – partial dissociation,
both ions and molecules
CH3COOH
NH3
Need to know
equilibrium
constant
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 =
14 - pOH
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 = 11.2681
• Calculate the pH of a 1 molar ethanoic
acid solution that is only 1.4% ionised
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
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)
Methyl orange
H+(aq) + In-(aq, colour 2)
•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
Draw rough trend graph
Name of
Indicator
Approx
Range
Methyl
Orange
3.1-4.4
Litmus
5-8
Acid
Colour
Lower pH
Base Colour
Higher pH
yellow
red
Phenolphth 8.3-10
alein
red
blue
colourless
pink
Acid Base Titration Curves
Strong Acid – Strong Base
Weak Acid – Strong Base
Strong Acid – Weak Base
Weak Acid – Weak Base
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
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
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.
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.
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
Page 261, 262
Question NB to practise