Acids and Bases in Aqueous Solution
Acid-Base Theory
An acid is a proton donor and
a base is a proton acceptor
(a) acids as proton donating species (molecule or ion)
e.g. sulfuric acid H2SO4
H+ +
_______________
b) bases as proton accepting species (molecule or ion)
e.g. ammonia NH3 NH3
+ H+
_________
(c) amphiprotic species i.e. a molecule or ion which can act as an acid or a base
hydrogen carbonate ion HCO3-
HCO3– + H+
_________
HCO3–
H+ + _________
(d) conjugate pairs
When a species behaves as an acid it donates a proton and forms a conjugate base
When a species behaves as a base it accepts a proton and forms a conjugate acid
eg:
NH3
+
____________
H2O
NH4+
+
____________
______________
OH –
_______________
An acid and its conjugate base differ
by a PROTON (hydrogen ion)
Compare the strength of H2SO4 and HSO4-
Strength of Acids and Bases
Strengths of acids may be compared by the extent of their proton transfer to
water forming aqueous hydronium ions. In other words, a strong acid ionises to
a very great extent.
strong acid: HCl + H2O
___________ + ___________
practically all HCl molecules have ionised
weak acid:
CH3COOH + H2O
___________ + ___________
few CH3COOH molecules have ionised.
Examples of strong acids: ______________________________________________
Examples of weak acids: _______________________________________________
The strengths of bases may be compared by the extent to which they produce OHions in solution.
1. Strong base:
NaOH
_________ + _________
This ionisation proceeds to a very great extent
2. Weak base:
NH3 + H2O
________ + ________
This ionisation proceeds to a very limited extent.
Examples of Strong Bases: ___________________________________________
Examples of Weak Bases: ____________________________________________
Acid-Base nature of Salts
A salt can be acidic, basic or neutral depending how the cation or anion present reacts
with water.
For example, ammonium (NH4+) salts tend to be acidic, because the NH4+ ion is the
conjugate acid of a weak base, ie: it is a strong acid. This can be shown by the
following hydrolysis reaction:
NH4+
Cl– +H2O
NH3 + H+
happens to a _________ extent
HCl + OH – happens to a ____________ extent
So ammonium chloride is an acidic salt, because the ammonium ion came form a
weak base (NH3) and the chloride came from a strong acid (HCl). Sodium ethanoate is
a basic salt because sodium came from a strong base (NaOH) and ethanoate came
from a weak acid (CH3COOH)
 Worked example 14
All of the hydrogen halides are acidic in aqueous solution. Explain any trend in acid
strength related to the halogen's position in the group.
Kw and pH Calculations
Water ionises to H+ ions and OH- ions to a very limited extent.
H2O
H+ + OH –
The equilibrium constant for this ionisation is very small at 10-14 (at 25C)
i.e.
K w  10
 14
[ H  ][OH  ]

1
[a useful extension of this is pKw = 14 = pOH + pH]
Because water is neutral, [H+] = [OH-] = ________
pH   log[ H  ]
 Worked example 15
Find the pH of a 0.001 molL-1 solution of Ca(OH)2 solution.
Amphoteric Substances
An amphoteric substance reacts with both acids and bases. Chromium, zinc and
aluminium are amphoteric metals and they form amphoteric oxides and
amphoteric hydroxides. These metals form complex ions with bases.
eg, Aluminium:
Al(s) + 2OH –(aq) + 6H2O
2Al(s) + 6H+(aq)
2[Al(OH)4] –(aq) + 3H2(g)
2Al3+(aq) + 3H2(g)
Complex Ions
A complex ion consists of a positive ion (usually a
transition metal ion) surrounded by a number of
ligands (ions or polar molecules).
Eg: [Fe(OH)6]3+, [Al(OH)4]3+ , [Cu(NH3)4]2+
Cu2+
Volumetric Analysis
You must know the steps that need to be followed when carrying out volumetric
analysis - they are not presented here, but are available in your text book and
laboratory manual. Be able to describe (i) how to make a primary standard, (ii) do an
accurate titration, (iii) explain these curves and (iv) discuss the sources of error.
pH
14
pH
14
0
0
Volume of acid added
Volume of acid added
14
pH
pH
14
0
0
Volume of acid added
Volume of acid added
Indicators:
Indicator
Methyl Orange
Phenolphthalein
Useful pH range
3.1 – 4.4
8.3 - 10
Colour Change
red - yellow
colourless - red (pink)
It is crucial that an indicator is chosen so that the end point of the titration
corresponds to the equivalence point of the neutralisation.
Acid-Base Calculations
Stoichiometry is applied in acid-base calculations in exactly the same way as it is
applied in any other area of Chemistry. The four steps of stoichiometry apply (see
earlier)
 Worked example 16:
Barium hydroxide is normally hydrated, this means it is of the form: Ba(OH)2.xH2O.
0.816 g of a sample of hydrated barium hydroxide was dissolved in 25.0 mL of water
and titrated against 0.177 molL-1 hydrochloric acid. The titration required exactly
29.3 mL of the acid for complete neutralisation.
(a)
State a suitable indicator for the titration. Explain your choice.
(b)
Find the correct formula for the Ba(OH)2.xH2O.
 Practice Questions
1.
Ammonium nitrate solution is
(a)
(b)
(c)
(d)
(e)
2.
Which of the following solutions has the lowest pH?
(a)
(b)
(c)
(d)
3.
0.1 M ethanoic acid
0.1 M hydrochloric acid
0.2 M sodium hydroxide
0.2 M nitric acid.
Which of the following descriptions best describes 10 mol L-1 ammonia?
(a)
(b)
(c)
(d)
4.
Acidic because nitrate ions react with water to give HNO3
Basic because ammonium ions react with water to give OHNeutral because ammonia is basic and nitric acid is acidic.
Acidic because ammonia reacts with water liberating H3O+ ions.
Basic because it contains ammonia which is a base.
a dilute solution of a weak base
a concentrated solution of a weak base
a dilute solution of a strong base
a strong solution of a weak base
Which of the following has a pH less than 7?
(a)
(b)
(c)
(d)
0.01 mol L-1 ammonium chloride
0.01 mol L-1 calcium hydroxide
0.01 mol L-1 potassium nitrate
0.01 mol L-1 sodium ethanoate
6.
Ethanoic acid is titrated against sodium hydroxide (in the burette), using
methyl orange as the indicator. Which of the following statements about the
titration is true?
(a)
(b)
(c)
(d)
7.
The end point and equivalence point occur at the same time.
The end point occurs after the equivalence point.
The end point occurs before the equivalence point.
No end point would occur.
Classify the following solutions as acidic, basic or neutral. Where the solution
is basic or acidic, write a hydrolysis equation to support your answer.
KH2PO4
8.
MgF
LiHCO3
KCl
Chemists can standardise sodium hydroxide solutions with the monprotic acid
potassium hydrogen phthalate (KHC8H4O4) as a primary standard. A 25.0 mL
aliquot of a sodium hydroxide solution requires 23.1 mL of 0.0994 mol L-1
potassium hydrogen phthalate for complete reaction.
(a)
Suggest an indicator for the titration, explain your choice.
(b)
Explain what is meant by a primary standard.
(c)
Why do you think the chemist standardised the sodium hydroxide
solution?
(d)
9.
Find the concentration of the sodium hydroxide solution.
Orange juice contains the weak acid citric acid. The
formula is shown in the figure.:
O
H OH H
O
C C C C C
HO
H C H
HO O
(a)
State what is meant by the term weak acid.
(b)
Each molecule of citric acid contains 8
hydrogen atoms. Only three of these can be considered as acidic
hydrogens. Identify these by circling them on the diagram.
(c)
24.2 g of a particular brand of orange juice was placed in a conical
flask. Titration with sodium hydroxide solution showed that the orange
juice contained 0.700 % citric acid by mass.
(d)
(i) Calculate the mass, and hence the number of moles of citric acid in
the sample of orange juice.
(ii) How many moles of sodium hydroxide were required to react
completely with this amount of citric acid.
OH
Exemplar exam (Section 1)
Question 15
Which of the following solutions can be added to 100.0 mL of a 0.2 mol L-1 NH3
solution to produce a buffer?
I 100.0 mL of 0.20 mol L-1 NH4Cl(aq)
II 50 mL of 0.20 mol L-1 HCl(aq)
III 100.0 mL of 0.20 mol L-1 HCl(aq)
a)
b)
c)
d)
I only
I and II only
II and III only
I, II and III
Question 16
Pure water undergoes self-ionisation according to the equation:
2H2O()
H3O+(aq) + OH-(aq)
The equilibrium constant for the reaction is:
1.0 x 10-14 at 25oC
5.5 x 10-13 at 100oC
Which one of the following statements is correct?
a) At 100oC the pH of water is less than 7, but the water is still neutral.
b) At 100oC the pH of water is less than 7, and therefore the water is acidic.
c) At 100oC the pH of pure water is greater than 7.0, and therefore the water
is basic.
d) At 100oC the pH of water is greater than 7, and therefore the water is
acidic.
Question 17
Potassium permanganate solution is standardised using the following procedure:
sulfuric acid is added to 25.00 mL of a standard solution of oxalic acid (ethandioic
acid) and the mixture warmed and titrated with potassium permanganate solution.
Which one of the following will cause an error?
a) Leaving the conical flask (titration vessel) wet after rinsing with distilled
water.
b) Rinsing the burette with distilled water, then with a little of the
permanganate solution, and then filling it with the permanganate solution.
c) Rinsing the pipette with distilled water, then using it to dispense the oxalic
acid solution.
d) Washing down the sides of the conical flask with distilled water from time
to time during the titration.
Buffers
Exemplar from (Section 2)
[9 marks]
The pH of blood is maintained through buffering. The major buffer system present in
blood is based on a carbonic acid/hydrogencarbonate ion buffer. When carbon
dioxide enters the blood stream, the following reaction occurs:
CO2(g) + 2 H2O()
H2 CO3( aq ) + H 2 O (  )
+ H3 O(aq)
HCO 3(aq)
The presence of carbonic acid and hydrogencarbonate ions maintains the pH of
blood to about 7.4.
(a)
Excess HCO 3(aq)
produced as a result of strenuous exercise is removed from
the blood by the kidneys.
(i)
Describe the immediate change in the pH of the blood.
[1 mark]
__________________________________________________________
__________________________________________________________
(ii)
Explain the immediate change in the pH.
[2 marks]
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________
(b)
Hyperventilation, or rapid breathing, decreases the amount of carbon dioxide in
the lungs. As a consequence the concentration of carbon dioxide dissolved in
the blood also decreases.
(i)
Describe the immediate change in the pH of the blood. [1 mark]
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
(ii)
Explain the immediate change in the pH. [2 marks]
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________
(iii)
Describe how the buffer system counteracts this change in pH. [3 marks]
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________
_______________________________________________________________