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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.
Ionisation
Ionic Product
pH
Logarithm
Kw
Indicator
pH scale
Strong/weak acids
Strong/Weak bases
pH Curve
End-Point
Dissociation Constant
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
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
Acid–Base Concentrations in Solutions
Acid–Base Concentrations in Solutions
OH-
H+OH-
OH-H+
H+
[H+] = [OH-] [H+] > [OH-] [H+] < [OH-]
acidicsolution
neutralsolution
basicsolution
co
nc
en
trat
ion
(m
ole
s/L
)
10-14
10-7
10-1
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.
Neutral Weak Alkali
Strong Alkali
Weak Acid
Strong Acid
7 8 9 10 11 12 133 4 5 62 141 7 8 9 10 11 12 133 4 5 62 141 9 10 11 123 4 5 621
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 pHUniversal Indicator Paper
Universal Indicator Solution
pH meter
Measuring pHMeasuring 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
Limitations of pH ScaleThe 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
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.
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+ ]
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 increasesA 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
•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
Acid – Base Concentrations and pH
pH = 3
pH = 7
pH = 11
OH-
H+OH-
OH-H+
H+
[H3O+] = [OH-] [H3O+] > [OH-] [H3O+] < [OH-]
acidicsolution
neutralsolution
basicsolution
co
nc
en
trat
ion
(m
ole
s/L
)
10-14
10-7
10-1
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
pH of Common Substances
Acidic Neutral Basic
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 MHousehold bleachHousehold ammonia
Lime waterMilk of magnesia
Borax
Baking sodaEgg white, seawaterHuman blood, tearsMilkSalivaRain
Black coffeeBananaTomatoesWineCola, vinegarLemon juice
Gastric juice
Mor
e ba
sic
Mor
e ac
idic
pH [H+] [OH-] pOH
7 1 x 10-7 1 x 10-7 7
Calculations and practiceCalculations and practice
pH = – log10[H+]
• You will need to memorize the following:
pOH = – log10[OH–]
[H+] = 10–pH
[OH–] = 10–pOH
pH + pOH = 14
pH Calculations
pH
pOH
[H+]
[OH-]
pH + pOH = 14
pH = -log10[H+]
[H+] = 10-pH
pOH = -log10[OH-]
[OH-] = 10-pOH
[H+] [OH-] = 1 x10-14
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 concentrationonly need to know the concentration.
Strong acids are so named because they react completely with water, leaving no undissociated molecules in solution.
pH ExercisespH Exercisesa) 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
c) pH of solution where [H +] is 7.2x10-8M
pH = – log10 [H+]= – log10 [7.2x10-8]= 7.14
(slightly basic)
monoproticmonoprotic
diproticdiprotic
HA(aq) H1+(aq) + A1-(aq)
0.3 M 0.3 M 0.3 M
pH = - log10 [H+]
pH = - log10[0.3M]
pH = 0.48e.g. HCl, HNO3
H2A(aq) 2 H1+(aq) + A2-(aq)
0.3 M 0.6 M 0.3 M
pH = - log10[H+]
pH = - log10[0.6M]
pH = 0.78e.g. H2SO4
pH = ?
pH = 4.6
pH = - log10 [H+]
4.6 = - log10 [H+]
- 4.6 = log10[H+]
- 4.6 = antilog [H+]
Given:
2nd log
10x
antilog
multiply both sides by -1
substitute pH value in equation
take antilog of both sides
determine the [hydrogen ion]
choose proper equation
[H+] = 2.51x10-5 M
You can check your answer by working backwards.
pH = - log10[H+]
pH = - log10[2.51x10-5 M]
pH = 4.6
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.
For Weak Acids
pH = -Log10
For Weak Bases
pOH = Log10
pH = 14 - pOH
Calculating pH - Calculating pH - weak acidsweak acids
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
A weak acid is one which only partially dissociates in aqueous solution
pH of solutions of weak concentrationsWeak 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
pH of solutions of weak concentrationsWeak 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
Theory of Acid Base IndicatorsAcid-base titration indicators are quite often weak acids.
For the indicator HInThe 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¯).
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+].
Theory of Acid Base IndicatorsApplying 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-]
Theory of Acid Base IndicatorsSummary In acidic solution
HIn(aq) H+(aq) + In¯(aq)
In alkaline solution
Theory of Acid Base IndicatorsAcid-base titration indicators are also often weak bases.
For the indicator MOHThe equilibrium can be simply expressed as
MOH(aq, colour 1) OH-(aq) + M+
(aq, colour 2)
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-].
Theory of Acid Base IndicatorsApplying 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+]
Acid Base Titration CurvesStrong Acid – Strong Base Strong Acid – Weak Base
Weak Acid – Strong Base
25 cm3 of 0.1 mol dm-3 acid is titrated with 0.1 mol dm-3 alkaline solution.
Weak Acid – Weak Base
Choice of Indicator for TitrationIndicator 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
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
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.
Methyl Orange is
used as indicator for
this titration
Indicators for Weak Acid Strong Base 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.
Phenolphthalein is
used as indicator for
this titration
Indicators for Weak Acid Weak Base Titration
Neither phenolphthalein
nor methyl orange have
completely change colour
in the vertical section on
the pH titration curve
No indicator suitable
for this titration
because no vertical
section
indicator pH range
litmus 5 - 8
methyl orange 3.1 - 4.4
phenolphthalein 8.3 - 10.0
Colour Changes and pH ranges
Methyl Orange
Phenolphthalein
Universal indicator components
Indicator Low pH color Transition pH range High pH color
Thymol blue (first transition) red 1.2–2.8 orange
Methyl Orange red 4.4–6.2 yellow
Bromothymol blue yellow 6.0–7.6 blue
Thymol blue (second transition) yellow 8.0–9.6 blue
Phenolphthalein colourless 8.3–10.0 purple
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