Ionic Equilibrium II : Buffers, Indicators and Solubility Product

1
Ionic Equilibrium II Buffers, Indicators and
Solubility Product
18.1 Buffer Solutions 18.2 Calculations
Involving Composition and pH of Buffer
Solutions 18.3 Acid-base Indicators 18.4 Acid-base
Titrations 18.5 Solubility Product
2
Buffer Solutions
3
18.1 Buffer Solutions (SB p.153)
Buffer Solutions
A buffer solution is a solution that tends to
resist changes in pH when a small amount of acid
or base is added to it.
4
18.1 Buffer Solutions (SB p.153)
Acidic Buffer Solutions

• Prepared by mixing a weak acid and a salt
  containing its conjugate base
• e.g. ethanoic acid and sodium ethanoate
• The solution is acidic
• Used to resist pH changes in an acidic medium

5
18.1 Buffer Solutions (SB p.153)
Basic Buffer Solutions

• Prepared by mixing a weak base and a salt
  containing its conjugate acid
• e.g. aqueous ammonia and ammonium chloride
• The solution is basic
• Used to resist pH changes in an basic medium

6
18.1 Buffer Solutions (SB p.153)
Working Principle of Acidic Buffer Solutions
e.g. a buffer solution of ethanoic acid and
sodium ethanoate

• Sodium ethanoate is completely ionized in water
• Gives a large amount of ethanoate ions
• CH3COONa(aq) ?? CH3COO(aq)
  Na(aq)

7
18.1 Buffer Solutions (SB p.154)
Working Principle of Acidic Buffer Solutions

• Ethanoic acid is only slightly ionized in water

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18.1 Buffer Solutions (SB p.154)
Working Principle of Acidic Buffer Solutions

• When a small amount of an acid is added to the
  system
• ? The position of the equilibrium shifts to the
  left
• ? The H3O(aq) ions added are removed by
  forming the partially ionized CH3COOH
• ? The pH of solution remains almost constant

9
18.1 Buffer Solutions (SB p.154)
Working Principle of Acidic Buffer Solutions

• When a small amount of a base is added to the
  system
• ? The additional OH-(aq) ions react with
  undissociated CH3COOH molecules to form
  CH3COO-(aq) ions and H2O molecules
• OH(aq) CH3COOH(aq) ?? CH3COO(aq)
  H2O(l)
• ? The pH of solution remains almost constant

10
18.1 Buffer Solutions (SB p.154)
Working Principle of Acidic Buffer Solutions
Diagram showing how an acidic buffer solution
works
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18.1 Buffer Solutions (SB p.155)
Calculation of the pH of an Acidic Buffer Solution
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18.1 Buffer Solutions (SB p.155)
Calculation of the pH of an Acidic Buffer Solution
Putting H3O(aq) as the subject of the equation,
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18.1 Buffer Solutions (SB p.155)
Calculation of the pH of an Acidic Buffer Solution
Taking negative logarithms on both sides
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18.1 Buffer Solutions (SB p.155)
Calculation of the pH of an Acidic Buffer Solution

• CH3COONa is completely ionized and CH3COOH is
  only slightly ionized in the solution
• ? Assume that the CH3COO ions come from
  CH3COONa(aq) only
• ? The contribution of CH3COO ions from CH3COOH
  is negligible

15
18.1 Buffer Solutions (SB p.155)
Calculation of the pH of an Acidic Buffer Solution
CH3COO(aq) CH3COONa(aq)
salt CH3COOH(aq) acid
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18.1 Buffer Solutions (SB p.155)
Calculation of the pH of an Acidic Buffer Solution

• To calculate the pH values of acidic buffer
  solutions
• Shows that the pH of an acidic buffer solution
  depends on the ratio of the concentrations of the
  acid and the salt

17
18.1 Buffer Solutions (SB p.156)
Working Principle of Basic Buffer Solutions
e.g. a buffer solution of aqueous ammonia and
ammonium chloride

• Ammonium chloride is completely ionized in water
• Gives a large amount of ammonium ions
• NH4Cl(aq) ?? NH4(aq) Cl-(aq)

18
18.1 Buffer Solutions (SB p.156)
Working Principle of Basic Buffer Solutions

• Aqueous ammonia is only slightly ionized in water

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18.1 Buffer Solutions (SB p.156)
Working Principle of Basic Buffer Solutions

• When a small amount of an acid is added to the
  system
• The H3O(aq) ions added react with NH3
  molecules to form NH4 ions
• The additional H3O(aq) ions are neutralized
• The pH of the system remains almost constant
• H3O(aq) NH3(aq) ?? NH4(aq) H2O(l)

20
18.1 Buffer Solutions (SB p.156)
Working Principle of Basic Buffer Solutions

• When a small amount of a base is added to the
  system
• The OH(aq) ions added are removed by forming
  ammonia molecules which only dissociate slightly
• The pH of the system remains almost constant
• OH(aq) NH4(aq) ?? NH3(aq) H2O(l)

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18.1 Buffer Solutions (SB p.156)
Working Principle of Basic Buffer Solutions
Diagram showing how a basic buffer solution works
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18.1 Buffer Solutions (SB p.157)
Calculation of the pH of a Basic Buffer Solution
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18.1 Buffer Solutions (SB p.157)
Calculation of the pH of a Basic Buffer Solution
Putting OH-(aq) as the subject of the equation,
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18.1 Buffer Solutions (SB p.157)
Calculation of the pH of a Basic Buffer Solution
Taking negative logarithms on both sides
25
18.1 Buffer Solutions (SB p.157)
Calculation of the pH of a Basic Buffer Solution
NH4(aq) NH4Cl(aq) salt NH3(aq)
base
26
Calculations Involving Composition and pH of
Buffer Solutions
27
18.2 Buffer Solutions (SB p.158)
Calculations Involving Composition and pH of
Buffer Solutions
For acidic buffer solutions
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18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.158)
Calculations Involving Composition and pH of
Buffer Solutions
For alkaline buffer solutions
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18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.158)
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Acid-base Indicators
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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators
An acid-base indicator is a substance which
changes colour as the pH of the solution in which
it is dissolved changes
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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators

• The indicators used in acid-base titrations are
  either weak acids or weak bases
• Their conjugate bases or conjugate acids are of
  colours different from those of the indicators
  themselves

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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators
When the indicator is added to water
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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators

• In an acidic medium where H3O(aq) ions is
  high
• ? The position of the above equilibrium shifts
  to the left
• ? The dominant species in the solution will be
  HIn(aq) ? Shows colour 1

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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators

• In an alkaline medium where H3O(aq) ions is
  low
• ? The position of the above equilibrium shifts
  to the right
• ? The dominant species in the solution will be
  In-(aq) ? Shows colour 2

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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators
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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators
The product of Kc and H2O(l) gives another
constant called the indicator dissociation
constant, KIn
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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators
Putting H3O(aq) as the subject of the equation,
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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators
Taking negative logarithms on both sides,
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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators

• The ratio of In(aq) to HIn(aq) is related to
  the pH of the solution (i.e. H3O(aq) of the
  solution)
• The colour of the indicator changes with
  H3O(aq)

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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators

• When H3O(aq) decreases
• ? In(aq) increases
• ? HIn(aq) becomes smaller

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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators

• When H3O(aq) increases
• ? In(aq) decreases
• ? HIn(aq) becomes larger

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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators
? Colour 1 of the indicator is observed
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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators
? Colour 2 of the indicator is observed
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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators

• When In-(aq) HIn(aq)
• ? A colour formed from the combination of
  colour 1 and colour 2 is observed

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18.3 Acid-base Indicators (SB p.164)
Acid-base Indicators

• The indicator changes from colour 1 to colour 2
  over a range of H3O(aq) from pH (pKIn 1)
  to pH (pKIn 1)
• The colour change of an acid-base indicator
  generally takes place over a range of two pH units

47
18.3 Acid-base Indicators (SB p.164)
Phenolphthalein

• Weak acid
• KIn value 7 1010 mol dm3
• When phenolphthalein dissolves in water
• ? its acidic form (HPh) is colourless
• ? its conjugate base (Ph) is pink

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18.3 Acid-base Indicators (SB p.164)
Phenolphthalein
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18.3 Acid-base Indicators (SB p.165)
Phenolphthalein
? The indicator appears colourless
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18.3 Acid-base Indicators (SB p.165)
Phenolphthalein
When Ph-(aq) HPh(aq)
? The indicator appears pale pink
pH pKIn log (7 1010) 9.15
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18.3 Acid-base Indicators (SB p.165)
Phenolphthalein
? The indicator appears pink
pH pKIn log 10 pKIn 1 log (7
1010) 1 10.15
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18.3 Acid-base Indicators (SB p.165)
Phenolphthalein
pH 9.15
pH 8.15
pH 10.15
Colours of phenolpthalein at pH 8.15, 9.15 and
10.15
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18.3 Acid-base Indicators (SB p.166)
Methyl Orange

• Weak base
• KIn value 2 104 mol dm3
• Methyl orange (Me) is yellow
• Its conjugate base (HMe) is red

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18.3 Acid-base Indicators (SB p.166)
Methyl Orange
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18.3 Acid-base Indicators (SB p.166)
Methyl Orange
? The indicator appears red
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18.3 Acid-base Indicators (SB p.166)
Methyl Orange
When Me(aq) HMe(aq)
? The indicator appears orange
pH pKIn log (2 104) 3.7
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18.3 Acid-base Indicators (SB p.166)
Methyl Orange
? The indicator appears yellow
pH pKIn log 10 pKIn 1 log (2
104) 1 4.7
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18.3 Acid-base Indicators (SB p.166)
Methyl Orange
pH 3.7
pH 2.7
pH 4.7
Colours of methyl orange at pH 2.7, 3.7 and 4.7
59
18.3 Acid-base Indicators (SB p.166)
Colour changes of some common indicators at
particular pH ranges
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Acid-base Titrations
61
18.4 Acid-base Titrations (SB p.168)
Acid-base Titrations
Titration is a method used in quantitative
analysis to determine the quantity or
concentration of a solution

• In acid-base titrations, an acid is added to an
  alkali or an alkali is added to an acid, until
  they just neutralize each other

62
18.4 Acid-base Titrations (SB p.168)
Acid-base Titrations
Equivalence point is the point at which
equivalent amounts of the acid and alkali have
reacted
The point at which the indicator changes colour
is called the end point
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18.4 Acid-base Titrations (SB p.168)
Titration Curves
A titration curve is a curve showing the changes
in pH in the process of a titration
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18.4 Acid-base Titrations (SB p.169)
Titration Curves
Strong acid against strong alkali
Strong acid against weak alkali
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18.4 Acid-base Titrations (SB p.169)
Titration Curves
Weak acid against strong alkali
Weak acid against weak alkali
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18.4 Acid-base Titrations (SB p.169)
Titration Curves

• An ideal indicator for a titration is the one
  whose end point matches the equivalence point of
  the titration

67
18.4 Acid-base Titrations (SB p.171)
Double Indicator Titration

• The reaction between sodium carbonate and a
  strong acid takes place in two stages
• The completion of the reaction at different
  stages requires different indicators
• Called double indicator titration

68
Solubility Product
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18.5 Solubility Product (SB p.172)
Meaning of Solubility Product
Consider the dissolution of a sparingly soluble
salt MX(s) (e.g. AgCl) in water to form M(aq)
and X(aq)
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18.5 Solubility Product (SB p.172)
Meaning of Solubility Product

• The amount of MX(s) present in the solution does
  not affect its solubility

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18.5 Solubility Product (SB p.172)
Meaning of Solubility Product

• Incorporating the MX(s) into Kc of the above
  equilibrium gives a new constant known as the
  solubility product, Ksp
• Ksp M(aq)eq X(aq)eq

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18.5 Solubility Product (SB p.172)
Meaning of Solubility Product
Generally, MaXb(s) ?? aMb(aq) bXa(aq) the
solubility product is given by Ksp
Mb(aq)aeq Xa(aq)beq
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18.5 Solubility Product (SB p.172)
Meaning of Solubility Product
BaSO4(s) ?? Ba2(aq) SO42(aq) Ksp
Ba2(aq)eq SO42(aq)eq Ksp of BaSO4 at 298 K
1 1010 mol2 dm6
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18.5 Solubility Product (SB p.172)
Meaning of Solubility Product
Ag2CO3(s) ?? 2Ag(aq) CO32(aq) Ksp
Ag(aq)2eq CO32eq Ksp of Ag2CO3 at 298 K
8 1012 mol3 dm9
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18.5 Solubility Product (SB p.173)
Comparison between Ionic Product and Solubility
Product

• M(aq) X(aq) lt Ksp at a given temperature
• ? the solution is unsaturated
• ? more MX(s) can be dissolved in it

76
18.5 Solubility Product (SB p.173)
Comparison between Ionic Product and Solubility
Product
2. M(aq) X(aq) Ksp at a given
temperature ? the solution is saturated ? the
system is at equilibrium
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18.5 Solubility Product (SB p.173)
Comparison between Ionic Product and Solubility
Product
3. M(aq) X(aq) gt Ksp at a given
temperature ? the solution is supersaturated ? p
recipitation will occur until M(aq) X(aq)
Ksp
78
18.5 Solubility Product (SB p.174)
Common Ion Effect

• Ksp of a salt MX(s) is always constant at a given
  temperature
• Independent of the concentrations of the M(aq)
  and X(aq) ions
• Addition of a common ion (i.e. M(aq) or X(aq)
  ions) will change the position of the equilibrium

79
18.5 Solubility Product (SB p.175)
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The END
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18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.158)
Back
Let's Think 1
What is the buffering capacity of a buffer
solution and what determines the buffering
capacity of a buffer solution?
Answer
The buffering capacity of a buffer solution
refers to the amount of acid or base that a
buffer solution allows to add into it without
changes in its pH. The buffering capacity depends
on the degree of ionization of the buffering
molecules in the buffer solution. The buffering
capacity is higher if there is a larger reservoir
of a conjugate acid-base pair in the buffer
solution.
82
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.158)
Example 18-2A
A buffer solution is made by adding 4.1 g of
sodium ethanoate to 1 dm3 of a 0.01 M solution of
ethanoic acid. Calculate the pH of the buffer
solution. (Given Ka of ethanoic acid at 298 K
1.74 105 mol dm3 molar mass of sodium
ethanoate (CH3COONa) 82 g mol1 assume there
is no volume change on mixing)
Answer
83
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.158)
Example 18-2A
Back
84
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.158)
Example 18-2B
(a) Calculate the change in pH when 1 cm3 of 0.25
M sodium hydroxide solution is added to (i) 25
cm3 of pure water (ii) 25 cm3 of 0.1 M ethanoic
acid (iii) 25 cm3 of 0.1 M ethanoic acid
containing 0.002 mole of sodium
ethanoate. (Given Ka of ethanoic acid at 298 K
1.74 105 mol dm3 Kw of water at 298 K 1
1014 mol2 dm6 )
Answer
85
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.159)
Example 18-2B
86
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.160)
Example 18-2B
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18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.160)
Example 18-2B
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18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.160)
Example 18-2B
89
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.160)
Example 18-2B
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18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.160)
Example 18-2B
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18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.160)
Back
Example 18-2B
Answer
(b) Comment on the results obtained.
(b) A drastic pH change (4.98 units) is shown
when 1 cm3 of 0.25 M sodium hydroxide solution is
added to 25 cm3 of pure water as there is no
buffering effect. There is also a significant
change in pH (0.93 unit) when 1 cm3 of 0.25 M
sodium hydroxide solution is added to 25 cm3 of
0.1 M ethanoic acid. This shows that ethanoic
acid solution is not a good buffer solution as
there is no sodium ethanoate present. However,
when the same volume of sodium hydroxide solution
is added to the buffer solution containing 25 cm3
of 0.1 M thanoic acid and 0.002 mole of sodium
ethanoate, the change in pH is very small (0.1
unit). This indicates clearly that a buffer
solution is very effective in resisting the
change in pH when a small amount of base is added
to it.
92
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.162)
Example 18-2C
How many grams of ammonium chloride would you add
to 100 cm3 of 0.1 M aqueous ammonia in order to
prepare a basic buffer solution of pH 9.0?
(Given Kb of ammonia at 298 K 1.74 105 mol
dm3 molar mass of ammonium chloride (NH4Cl)
53.5 g mol1)
Answer
93
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.162)
Back
Example 18-2C
94
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.162)
Example 18-2D
Answer
95
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.162)
Example 18-2D
96
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.162)
Example 18-2D
Answer
97
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.162)
Example 18-2D
Back
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18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.163)
Check Point 18-2

• Calculate the pH of a buffer solution prepared by
  mixing 30 cm3 of 0.1 M ethanoic acid and 20 cm3
  of 0.08 M sodium ethanoate.
• (Given Ka of ethanoic acid at 298 K 1.74
  105 mol dm3)

Answer
99
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.163)
Check Point 18-2
100
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.163)
Check Point 18-2
(b) A buffer solution contains 0.100 mole of
ethanoic acid and 0.025 mole of sodium ethanoate
dissolved in distilled water and made up to a
1.00 dm3 of solution. (i) Calculate the pH of
the buffer solution. (Given Ka of ethanoic acid
at 298 K 1.74 105 mol dm3)
Answer
101
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.163)
Check Point 18-2
102
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.163)
Check Point 18-2
Check Point 18-2
(b) A buffer solution contains 0.100 mole of
ethanoic acid and 0.025 mole of sodium ethanoate
dissolved in distilled water and made up to a
1.00 dm3 of solution. (ii) When 103 mole of
HCl(aq) is added to the buffer solution, what is
the pH of the solution? Assume there is no
volume change. (Given Ka of ethanoic acid at
298 K 1.74 105 mol dm3)
Answer
103
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.163)
Check Point 18-2
104
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.163)
Check Point 18-2
(b) A buffer solution contains 0.100 mole of
ethanoic acid and 0.025 mole of sodium ethanoate
dissolved in distilled water and made up to a
1.00 dm3 of solution. (iii) When 102 mole of
HCl(aq) is added to the buffer solution, what is
the pH of the solution? Assume there is no
volume change. (Given Ka of ethanoic acid at
298 K 1.74 105 mol dm3)
Answer
105
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.163)
Check Point 18-2
106
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.163)
Check Point 18-2
(b) A buffer solution contains 0.100 mole of
ethanoic acid and 0.025 mole of sodium ethanoate
dissolved in distilled water and made up to a
1.00 dm3 of solution. (iv) When 5.0 102 mole
of NaOH(aq) is added to the buffer solution,
what is the pH of the solution? Assume there is
no volume change. (Given Ka of ethanoic acid
at 298 K 1.74 105 mol dm3)
Answer
107
18.2 Calculations Involving Composition and pH
of Buffer Solutions (SB p.163)
Check Point 18-2
Back
108
18.3 Acid-base Indicators (SB p.166)
Example 18-3
Answer
109
18.3 Acid-base Indicators (SB p.166)
Example 18-3
110
18.3 Acid-base Indicators (SB p.166)
Example 18-3
Answer
111
18.3 Acid-base Indicators (SB p.166)
Example 18-3
(b) Since the method is based on the visual
matching of colours, it is subjected to human
errors. In addition, the buffer solution used for
the experiment must have a pKa value very close
to the pKIn value of the indicator, and the
difference should not exceed one unit. Otherwise,
the colour of the indicator in the buffer
solution could never match that of the indicator
solution prepared.
Back
112
18.4 Acid-base Titrations (SB p.168)
Let's Think 2
What do you understand by the terms equivalence
pointand end point of a titration?
Answer
The point at which the acid and alkali just
neutralize each other is called the equivalence
point of the reaction. The point at which the
indicator changes colour is called the end point
of the titration which indicates the completion
of the reaction.
Back
113
18.4 Acid-base Titrations (SB p.169)
Let's Think 3
There is no suitable indicator to detect the
equivalence point of the titration of a weak acid
against a weak alkali. Do you know why?
Answer
In the titration of a weak acid against a weak
alkali, there is not a sharp change in pH at the
equivalence point. Therefore, there is no
suitable indicator that can indicate the
equivalence point of the titration.
Back
114
18.4 Acid-base Titrations (SB p.170)
Example 18-4
When 0.1 M hydrochloric acid was titrated against
25.0 cm3 of a solution containing sodium
hydrogencarbonate and sodium carbonate, 11.20 cm3
of the acid was used to decolourize
phenolphthalein. Upon reaching this end point,
methyl orange was added. A further 28.80 cm3 of
the acid was needed to turn methyl orange to
orange colour. Calculate the concentrations of
sodium carbonate and sodium hydrogencarbonate in
the original solution.
Answer
115
18.4 Acid-base Titrations (SB p.170)
Example 18-4
NaHCO3(aq) reacts with HCl(aq) in one
stage. NaHCO3(aq) HCl(aq) ?? NaCl(aq) H2O(l)
CO2(g) .............. (1) The colour change
of methyl orange from yellow to orange indicates
that the end point of the titration is
reached. For Na2CO3(aq), the reaction takes place
in two stages. Stage 1 Na2CO3(aq) HCl(aq) ??
NaHCO3(aq) NaCl(aq) ..... (2) Stage 2
NaHCO3(aq) HCl(aq) ?? NaCl(aq) H2O(l)
CO2(g) ... (3) The colour change of
phenolphthalein indicates that the end point of
reaction (2) is reached. The colour change of
methyl orange indicates that the end point of
reaction (3) is reached. The amounts of HCl(aq)
required for reaction (2) and reaction (3) are
the same. Therefore, 11.20 cm3 of 0.1 M HCl(aq)
was used to convert Na2CO3(aq) to NaHCO3(aq) only.
116
18.4 Acid-base Titrations (SB p.170)
Example 18-4
117
18.4 Acid-base Titrations (SB p.170)
Example 18-4
Back
118
18.4 Acid-base Titrations (SB p.171)
Check Point 18-4

• Name a suitable indicator for the titration of
• (i) hydrochloric acid against aqueous ammonia
• (ii) ethanoic acid against potassium hydroxide
  solution
• (iii) nitric acid against sodium hydroxide
  solution.

Answer
(a) (i) Methyl orange (ii) Phenolphthalein (iii)
Methyl orange or phenolphthalein
119
18.4 Acid-base Titrations (SB p.171)
Check Point 18-4
(b) Sketch the change in pH that occurs when
hydrochloric acid is added to sodium carbonate
until it is in excess. Explain your answer.
Answer
120
18.4 Acid-base Titrations (SB p.171)
Check Point 18-4
121
18.4 Acid-base Titrations (SB p.171)
Check Point 18-4
The reaction between hydrochloric acid and sodium
carbonate solution takes place in two stages. The
first stage involves the reaction of hydrochloric
acid and sodium carbonate solution to form sodium
chloride and sodium hydrogencarbonate. The
equivalence point of this reaction is at about pH
8. This point can be determined quite accurately
with the use of phenolphthalein. Stage 1
Na2CO3(aq) HCl(aq) ?? NaHCO3(aq) NaCl(aq) The
sodium hydrogencarbonate formed at stage 1 is
able to react with hydrochloric acid to form
sodium chloride, carbon dioxide and water. The
equivalence point of this reaction is at about pH
4, and this can be determined accurately with the
use of methyl orange. Stage 2 NaHCO3(aq)
HCl(aq) ?? NaCl(aq) CO2(g) H2O(l )
122
18.4 Acid-base Titrations (SB p.171)
Check Point 18-4
(c) 25 cm3 of a solution containing sodium
hydroxide and sodium carbonate was titrated with
0.05 M hydrochloric acid and phenolphthalein was
used as the indicator. After 18.50 cm3 of the
acid was added, phenolphthalein turned
colourless. Methyl orange was then added. As the
titration continued, a further 10.0 cm3 of acid
was needed to turn methyl orange to red colour.
Calculate the concentration of sodium hydroxide
and sodium carbonate in the original solution.
Answer
123
18.4 Acid-base Titrations (SB p.171)
Check Point 18-4
(c) Sodium hydroxide reacts with hydrochloric
acid in one stage only. NaOH(aq) HCl(aq) ??
NaCl(aq) H2O(l ) ....... (1) The
colour change of phenolphthalein indicates the
end point of reaction (1). However, sodium
carbonate solution reacts with hydrochloric acid
in two stages. Stage 1 Na2CO3(aq) HCl(aq) ??
NaHCO3(aq) NaCl(aq) .. (2) Stage 2
NaHCO3(aq) HCl(aq) ?? NaCl(aq) H2O(l )
CO2(g) .. (3) The colour change of
phenolphthalein indicates the end point of
reaction (2). The colour change of methyl orange
indicates the end point of reaction (3).
124
18.4 Acid-base Titrations (SB p.171)
Check Point 18-4
125
18.4 Acid-base Titrations (SB p.171)
Check Point 18-4
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126
18.5 Solubility Product (SB p.172)
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Let's Think 4
Are compounds having a higher solubility product
more soluble than those having a smaller
solubility product?
Answer
No. Although the solubility product is related to
the molar solubility of a salt, different salts
have different stoichiometries. Therefore, their
solubilities cannot be compared by their Ksp
values only.
127
18.5 Solubility Product (SB p.172)
Example 18-5A
The solubility of lead chromate(VI) (PbCrO4) is
4.5 105 g L1. Calculate the solubility
product of the compound.
Answer
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128
18.5 Solubility Product (SB p.172)
Example 18-5B
The solubility product, Ksp, of silver bromide
(AgBr) is 7.7 1013 mol2 dm6. Calculate the
solubility of the salt in mol dm3.
Answer
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129
18.5 Solubility Product (SB p.174)
Example 18-5C
Is a precipitate expected to form at equilibrium
when 50.0 cm3 of 0.001 0 M BaCl2(aq) is added to
50.0 cm3 of 0.000 10 M Na2SO4(aq)? (Given Ksp
for barium sulphate 1.1 1010 mol2 dm6.
Assume that the total volume of solution, after
mixing, equals the sum of the volumes of the
separate solutions.)
Answer
130
18.5 Solubility Product (SB p.174)
Example 18-5C
131
18.5 Solubility Product (SB p.174)
Example 18-5C
Ba2(aq) SO42(aq) (5.0 104) (5.0
105) 2.5
108 mol2 dm6 Since Ba2(aq) SO42(aq) (2.5
108 mol2 dm6) is greater than the Ksp (1.1
1010 mol2 dm6), barium sulphate precipitate is
expected to be formed.
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132
18.5 Solubility Product (SB p.175)
Example 18-5D
Calculate the solubility of silver chloride
in (a) pure water (Given Ksp of silver chloride
at 298 K 1.7 1010 mol2 dm6)
Answer
133
18.5 Solubility Product (SB p.175)
Example 18-5D
134
18.5 Solubility Product (SB p.175)
Example 18-5D
Calculate the solubility of silver chloride
in (b) 0.01 M sodium chloride solution (Given
Ksp of silver chloride at 298 K 1.7 1010
mol2 dm6)
Answer
135
18.5 Solubility Product (SB p.175)
Example 18-5D
136
18.5 Solubility Product (SB p.175)
Example 18-5D
137
18.5 Solubility Product (SB p.175)
Example 18-5D
Calculate the solubility of silver chloride
in (c) 0.5 M sodium chloride solution (Given Ksp
of silver chloride at 298 K 1.7 1010 mol2
dm6)
Answer
138
18.5 Solubility Product (SB p.175)
Example 18-5D
139
18.5 Solubility Product (SB p.175)
Example 18-5D
? The solubility of solid AgCl is 8.65 1011
mol dm3. The solubility of silver chloride
decreases further due to the presence of large
amounts of dissolved chloride ions.
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140
18.5 Solubility Product (SB p.176)
Check Point 18-5

• Bismuth sulphide (Bi2S3) has a solubility of 1.0
  1015 mol dm3 at 25 oC. Calculate the Ksp of
  bismuth sulphide.

Answer
141
18.5 Solubility Product (SB p.176)
Check Point 18-5
142
18.5 Solubility Product (SB p.176)
Check Point 18-5
(b) The solubility product of silver phosphate,
Ag3PO4, is 3.4 1014 mol4 dm12 at 298 K.
Calculate its solubility in mol dm3 at 298 K.
Answer
143
18.5 Solubility Product (SB p.176)
Check Point 18-5
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