Topic 8: ACIDS and BASES
8. 1. Theories of acids and bases
8.1.1 Define acids and bases according to the Brønsted–Lowry and Lewis theories.
8.1.2 Deduce whether or not a species could act as a Brønsted–Lowry and/or a Lewis acid or base.
8.1.3 Deduce the formula of the conjugate acid (or base) of any Brønsted–Lowry base (or acid).
Bronsted-Lowry theory (applies in aqueous solutions)
A Bronsted-Lowry acid is a proton donor
A Bronsted-Lowry base is a proton acceptor

Bronsted-Lowry acids must have a hydrogen/proton which can be released easily as an ion (a
replaceable hydrogen). This happens if the hydrogen is part of a weak bond or a polar bond. If the
acid has 1 replaceable hydrogen it is called a monoprotic acid.

The hydrogen ion (or proton as it is a particle with only a proton) released does not exist on its own
in a solution as it is always accepted by another specie i.e. a base e.g. water.
HA (l)
+
(acid molecule)
The above equation can also be written as:


H2O (l)
(base)
HA (l)
H3O+ (aq)

H+ (aq)
+
+
A- (aq)
A- (aq)
When acid molecules break up into ions in water we call it dissociation. Hydrogen ions are often
called protons because they are H atoms, which have lost an electron.
Examples of equations showing the dissociation/ionization of acid molecules.
HCl
+
H2O

HNO3
+
H2O

+
H2SO4
+
H2O

+
CH3COOH
+
H2O

+
H3O+ /H+(aq)
+
Cl-(aq)
Other molecules are said to be acidic because they react with water to form an acid. e.g sulfur trioxide
is called an acidic oxide because of the reaction.
SO3
+
H2O

H2SO4
Most non metal oxides are also acids because they act like acids i.e. they react with bases to produce
a salt and water (neutralization reaction)

Bronsted-Lowry bases are either

substances which contain ions like OH-, O2- , HCO32- or CO32-. These ions can accept the
proton/H+ (s) from the acid (because they contain atoms with 1 or more lone pairs);
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OH- (aq) +
CO32- (s) +
2H+ (aq)  H2O (l) + CO2 (g)
O2- (s) +
 H2O (l)
2H+ (aq)
H+ (aq)  H2O (l) + CO2 (g)
HCO3- (s) +

 H2O (l)
H+ (aq)
molecules that have lone pairs (e.g. ammonia and water); during an acid-base reaction, the
base forms a co-ordinate sigma bond with the proton from the acid e.g. when ammonia and
hydrogen chloride react to form ammonium chloride:
NH3
+
 NH4+ + Cl-
HCl
In this reaction the acid donates a proton to the base making the base a cation and the
acid an anion; attraction between the two of them results in an ionic bond and the formation
of NH4Cl.
SO WHEN A BASE ACCEPTS A PROTON IT FORMS A DATIVE BOND WITH IT!!!
According to the Bronsted-Lowry theory any reaction which involves the
transfer of protons from one species to another is an acid-base reaction.
Some substances can act as both and they are called amphoteric; what they behave like (either acid or
base) depends on the other chemical in the reaction. For instance if the other specie is a stronger acid
than our specie in question acts as a base, if the other specie is a stronger base than it will act as an
acid.
Conjugate acid-base pairs
All Bronsted-Lowry acid-base reactions like:

the ionisation of acid molecules
HCl (g)
+
H2O (l)
we can also write this as

H3O+ (aq)
+

HCl (aq)
Cl- (aq)
H+ (aq)
+
Cl- (aq)
the ionisation of alkali molecules
NH3 (g)
+
we can also write this as


acid-base reactions
H2O (l)

NH4+ (aq)
NH4OH (aq)
HCl (aq) +

NH3 (aq)
+
OH- (aq)
NH4 + (aq)

+
NH4+ (aq)
OH- (aq)
+
Cl- (aq)
All the above reactions are reversible processes (which means an equilibrium exists).
This is the case even when it involves strong acid or bases like hydrochloric acid and sodium
hydroxide. In the case of strong acids and bases the reversible reaction is ignored.
Within a reversible system, just like the forward reaction, the reverse reaction can also be considered a
Bronsted-Lowry acid-base reaction as it also involves a transfer of protons but in the opposite direction.
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In the equilibrium process there are then two acids and two bases; an acid and a base on each side of
the equation.
These acids and bases in such an equilibrium process are related: the acid on the reactant side is
related to the base (this base is called the conjugate base of the acid) on the product side in the
following way:
An acid + its conjugate base (= acid - H+ )
=
conjugate acid-base pair
A base + its conjugate acid (= base + H+)
=
conjugate acid-base pair
An acid dissociates into a proton and its conjugate base which is its anion.
The base on the reactant side is related in the same way to the acid on the product side; this acid is
called the conjugate acid of the base.
Example:
HCl (aq) +
acid
NH3 (aq)

NH4+ (aq)
conjugate acid
of the base NH3
Conjugate acid-base pairs from the above example:


Cl- (aq)
+
base
conjugate base
of the acid HCl
HCl and ClNH3 and NH4+ .
A conjugate acid-base pair is a pair of species which differ by the presence of one H+ only.
When an acid donates a proton it becomes a conjugate base. When a base accepts a proton it
becomes a conjugate acid.
A conjugate base behaves like a base while a conjugate acid behaves like an acid in the reverse
reaction.
As stated earlier some substances can behave like both acids and bases depending on the reaction
they are involved in; they are referred to as amphotheric substances eg water, HSO4- and NH4+ .
Exercises
1. In which one of the following reactions is the species in bold type behaving as a base?
A. 2NO + O2  2NO2
C. NH4+ + H2O  NH3 + H3O+
B. CO32- + H+  HCO3D. Cu2+ + 2OH-  Cu(OH)2
2. Which one of the following is the conjugate base of the hydrogen sulfite ion, HSO 3- ?
A. H2SO3
B. H2SO3-
C. SO32 -
D. SO3-
3. What is the conjugate acid of: CH3COO-, HSO4-, NH3 , OH- , F- ?
4. What is the conjugate base of: HCl, H3O+ , HSO4- , NH3 ?
5. Which one of the following species, many of which are unstable, would you expect to be capable of
acting as a base?
A. CH4
B. CH3
C. CH3+
D. CH3-
6. Which one can function as both an acid and a base according to the Bronsted-Lowry definition?
A. HSIB chemistry topic 8
B. S2-
C. NH4+
D. Al3+
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7. According to the equations below, what is the conjugate acid of HPO42H3PO4 + H2O  H2PO4- + H3O+
H2PO4- + H2O  HPO42- + H3O+
HPO42- + H2O  PO43- + H3O+
A. H3PO4
B. H2PO4-
C. HPO42-
D. PO43-
8. All the following are conjugate acid-base pairs EXCEPT
A. OH- , O2-
B. H2CO3, HCO3-
C. HNO2, NO2+
D. HSO4- , SO42-
9. Which of the following is NOT an acid-base pair?
A. HNO3 / NO3-
B. NH3 / NH2-
C. H2SO4 / HSO4-
D. H3O+ /OH-
10. In each of the following equations identify the conjugate acid-base pairs
(a) HNO2 + H2O  H3O+ + NO2(b) SO42- + H3O+  H2O + HSO4(c) H- + H2O  H2 + OH(d) H3O+ + OH-  2H2O
(e) HSO3- + H2O  H2SO3 + OH(f) H2 C2O4 + H2O  HC2O4- + H3O+
Structures of conjugate acid-base pairs
Members of a conjugate acid-base pair always differ by a single proton. To recognise a conjugate acid
or base more quickly, their written structures should always make clear the approximate location of the
proton that is transferred.
This applies especially to organic conjugate acid-base pairs.
Examples:
 ethanoic acid: CH3COOH and CH3COO ethanol: C2H5OH and C2H5O phenol: C6H6OH and C6H6ORelationship between acid and conjugate base and base and conjugate acid
 If an acid is strong its conjugate base is weak; the stronger the acid the weaker its conjugate base;
the more the equilibrium lies towards the products.
 The weaker the acid, the stronger the conjugate base which means the reverse reaction is favoured
and the equilibrium lies towards the reactants.
 The stronger the base the weaker its conjugate acid; the weaker the base, the stronger its conjugate
acid.
Lewis Theory: a much broader theory
Using this theory, a lot more reactions can now be described as ‘acid-base’ reactions.
All Bronsted/Lowry acid-base reactions are also Lewis
acid-base reactions!
Lewis acids: electron pair acceptors
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 A Lewis acid is a reactant which during a reaction accepts an electron pair from another substance;
 common Lewis acids are:
 boron and aluminium compounds in which neither element obeys the octet rule but has
empty orbitals in which it can place an electron pair;
 positive ions like the H+ which never exists on its own and either reacts with water to form
H3O+ or with a another base to form water. Other examples include CH3+ and Br+ (Hal+)
which you will study in organic chemistry.
 molecules containing positive centres (mostly organic molecules) as a result of polar bonds
within the molecule e.g. C in CO2 or in halogenalkanes and S in SO2.
 All Bronsted-Lowry acids are Lewis acids but not vice versa;
 Although strictly speaking dissociation needs to occur first before a Bronsted acid like HCl can act
as a Lewis acid; it is only the H+ which can accept the electron pair and therefore behave like a
Lewis acid.
 BF3 can act as a Lewis acid but not a Bronsted-Lowry.
 Lewis acids are electron deficient and are looking to gain electrons.
Lewis bases: electron pair donors
 A Lewis base is a reactant which during a reaction donates an electron pair; common Lewis bases
are:
 nitrogen compounds (e.g. ammonia),
 water and the hydroxide ion as they have a great tendency, because of their lone pairs, to donate
these electron pairs to Lewis acids.
 Also all Bronsted-Lowry bases are also Lewis bases because the Bronsted bases can only accept
protons by donating an electron pair to them.
 HIGHER LEVEL: Ligands (neutral molecules or anions with lone pairs) are Lewis bases that attach
themselves to metal ions which act like Lewis acids eg hydration of transition metal ions.
CuSO4.5H2O is in reality [Cu(H2O)4]2+ SO42- .H2O;
Lewis acid-base reactions
 A Lewis acid-base reaction involves the formation of a dative bond; a covalent bond in which both
electrons (lone pairs) are donated by the base to the acid; examples:
 H+ +
OH H2O
 NH3 + BF3  NH3BF3
(the boron in the above molecule used the unhybridised orbital to accommodate the dative bond)
Exercises
1. For each of the following species, state whether it is most likely to behave as a Lewis acid or Lewis
base. Explain your answer.
a) PH3
b) BCl3
c) H2S
e) Cu2+
d) SF4
2. All of the following species can function as Lewis bases EXCEPT
A. OHIB chemistry topic 8
B. N3-
C. NH3
D. NH4+
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3. Classify each of the following species as a Lewis acid or base.
(a) CO2
(c) I-
(b) H2O
(d) SO2
(e) NH3
(f) OH-
(g) BCl3
(h) NH4+
(i) H+
(j) N3-
4. Identify the Lewis acid and base in the following reactions.
(a) Hg2+ + 4CN-  Hg(CN)42(b) SnCl4 + 2Cl-  SnCl62(c) Co3+ + 6NH3  Co(NH3)63+
(d) CH3COO- + 2HF  CH3COOH
(e) CN- + H2O  HCN + OH(f) SO2 + H2O  H2SO4
(g) HNO2 + OH-  NO2- + H2O
+ HF2-
8. 2. Properties of acids and bases
8.2.1 Outline the characteristic properties of acids and bases in aqueous solution.
Acids
For an acid to show its acidic properties it must react with a base e.g. water.
During that reaction the acid molecule which has a covalently bonded (weak covalent bond or polar
covalent bond) hydrogen atom in it, splits into two ions: H+ and a negative non-metal ion like eg Cl- or
NO3- according to the following general equation:
HA
+
H2O

H3O+
+
A-
HA 
or
H+
+
A-
This reaction is a disscociation or ionisation reaction.
Within the scope of the syllabus we will limit our study to monoprotic acids which are acids with 1
hydrogen that they can donate. Examples of diprotic acids are sulfuric acid and carbonic acid.
It is the H3O+/H+ ions which give acids their properties:
 effect of acids on indicators:
* blue litmus turns red;
* phenolphthalein remains colourless;

* methyl orange: red-orange for low pH;
* universal indicator: deep red to red orange; pH  7;
react with metal oxides and hydroxides to form salts and water (=neutralisation); examples:

acid + metal oxide:

acid + metal oxide:

acid + metal hydroxide:
CuO(s) +
2HCl (aq)  CuCl2 (aq) +
H2O (l)
MgO(s) + 2CH3COOH (aq)  Mg(CH3COO)2 (aq) +
2NaOH(s) + H2SO4 (aq)  Na2SO4 (aq) +
H2O (l)
2H2O (l)
the above neutralisation reaction can be represented by the net ionic equation:
2OH- (aq) +
2H+ (aq)

2H2O (l)
as the sodium and sulfate ions are spectator ions.

react with the more reactive metals to produce a salt and hydrogen;
2HNO3 (aq) +
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Mg (s)
 Mg(NO3)2 (aq) + H2 (g)
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
react with metal carbonates and metal hydrogen carbonates to produce a salt, carbon dioxide and
water; examples:

CaCO3 (s) +

2NaHCO3 (s) +
2HCl (aq)
 CaCl2 (aq) +
 Na2SO4 (aq) +
H2SO4 (aq)
2H2O (l) + CO2 (g)
CO32- (s) + 2H+ (aq)  CO2 (g) + 2H2O (l))
(net ionic equation for the first one =

H2O (l) + CO2 (g)
react with ammonia to form ammonium salts, e.g.:
NH3 (s) +

HCl (aq)
 NH4Cl (s)
or
2NH3 (s) + H2SO4 (aq)  (NH4)2 SO4 (aq)
electrolytes as they ionise as a result of the dissolution caused by a polar solvent which in most
cases will be water; the greater the ion concentration the greater their conductivity.
Bases
 metal oxides, metal hydroxides, metal carbonates, metal hydrogen carbonates, amines;

ALKALIS = bases which are soluble in water; they form the OH- ion when they dissolve in
water (they either have the hydroxide ion as part of the compound and just dissociate – e.g. NaOH
– or they produce it when they react with water e.g. Li2O, NH3 or CO32-) .
 Na+ (aq) + OH- (aq)
NaOH (aq)
Li2O (s) + H2O (l)
 NH4 + (aq) + OH- (aq)
NH3 (g) + H2O (l)
CO32- + H2O (l)
effects on indicators:



red litmus turns blue;
phenolphthalein: pink;
 2Li+ (aq) + 2OH- (aq)
 HCO3- (aq) + OH- (aq)


methyl orange: yellow;
alkalis: universal indicator: green to deep violet; pH  7;

react with and neutralise acids to form salts and water;

alkalis are electrolytes as they dissociate as a result of the dissolution caused by a polar solvent
like water.
Exercise: For each of the following chemical reactions write a balanced formulae equation and an ionic
equation.
1.
2.
3.
4.
5.
6.
7.
magnesium and sulfuric acid
sodium hydrogen carbonate and hydrochloric acid
sulfuric acid and sodium chloride
calcium hydoxide and hydrochloric acid
copper and nitric acid
lead carbonate and nitric acid
iron with dilute sulphuric acid
8. 3. Strong and weak Bronsted acids and bases
8.3.1 Distinguish between strong and weak acids and bases in terms of the extent of dissociation, reaction with
water and electrical conductivity.
8.3.2. State whether a given acid or base is strong or weak.
8.3.3 Distinguish between strong and weak acids and bases, and determine the relative strengths of acids and
bases, using experimental data.
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Strength of an acid
Strong acids ionise (almost) or dissociate completely (equilibrium process which goes to virtual
completion) when they react with water. Strong acid = high concentration hydrogen ions in solution.
In weak acids not all of the acid molecules ionise, in fact most acid molecules remain molecules; the
weaker the acid the fewer the number of acid molecules that do so.
The weaker the acid the more the equilibrium of ionisation lies towards the reactants; the lower the
concentration of hydrogen ions in solution.
When a strong acid is added to water the ionisation is not really represented by an equilibrium:
e.g.:
HNO3

H+
+
NO3-
!!!!! We assume that all the H+ in the solution come from the acid and not the water which also
dissociates but only very slightly.
The strength of an acid is controlled by:

how easy it is to break the bond joining the hydrogen to the rest of the molecule; if the bond is
weak, the acid is strong. For example, although a covalently bonded molecule and very polar, HF
is a weak acid because the covalent bond is strong as a result of the small size of both the H and F
atoms. HI is a stronger acid than HBr and HCl as its hydrogen-halogen bond is the weakest.

how stable the anion is that is formed when the acid molecule donates its proton eg NO3-. If the
anion is stable the acid is strong as the anion does not readily react with a proton to form the acid
molecule again - it is a weak conjugate base;

the nature of the parent molecule eg ethanol is a weaker acid than ethanoic acid (more of this in
organics)
Therefore, a strong acid ionises readily or gives up very readily its hydrogen (either to water when it
dissolves or to a base).
Because of their different extent of ionisation, weak acids have a lower concentration of H+ in their
aqueous solutions, and therefore a higher pH, than strong acids of the same concentration.
Strong acids: HCl, HNO3, H2SO4 , HBr, HI, H3PO4.
Weak acids: CH3COOH (=ethanoic acid), H2CO3 (carbonic acid), HCOOH, citric acid, all
carboxylic/organic acids.
Exercise: For each of the above acids write an equation showing their dissociation in water.
Strength of bases
Strong bases dissociate (almost) completely (i.e. their ions separate) when they dissolve in water while
weak bases do not dissociate completely. A strong base accepts a proton very readily.
e.g.:
NaOH (s) 
OH- (aq)
+
Na+ (aq)
Weak bases have a lower concentration of OH- and therefore a lower pH than strong bases of the
same concentration. Most of their particles/molecules have not accepted protons.
Strong bases: all group 1 hydroxides and Ba(OH) 2 (all hydroxide ions in the solution are considered to
be coming from the base and not water);
Ba(OH) 2 (s)
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
2OH- (aq)
+
Ba2+ (aq)
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Weak bases: NH3 and amines.
NH3 (g) +
H2O (l)

NH4+ (aq)
+
OH- (aq)
Difference between strength and concentration.

the concentration of an acid or base refers to the number of acid or alkali molecules that are
present in 1L of the solution; when an acid is diluted more water is added so that there are fewer
acid molecules per 1L of solution;

strength refers to the degree of ionisation or dissociation that the acid or base undergoes when
dissolved in water.
Example: a hydrochloric acid solution and an ethanoic acid solution of the same concentration (eg
0.1M) will have different hydrogen concentrations - and therefore different pH’s - because of their
different strengths: the pH of the hydrochloric acid will be lower than the pH of the ethanoic acid
solution.
Some practical procedures to distinguish between strong and weak acids and bases.
Simple procedures to determine relative acidities or basicity of substances: It is important that in all
these tests, the acids or alkalis which are being compared are of the same
concentration.
All measurements taken in the procedures below are dependent on the amount of H+ ions present in
the solution. For example, when comparing two 1 mol dm-3 solutions of H2SO4 and HCl, there are
twice as many H+ ions in HsSO4 as in HCl so therefore the H2SO4 conducts better.

comparison of pH of substances; compare pH with known strong acids; if the pH is higher the acid
is weak.

comparison of conductivity of same concentrations; measure the conductivity of solutions of the
same concentration. The acidic or basic solutions with the largest concentration of ions will have
the highest conductivity as those with the highest conductivity have dissociated more.

compare rate of reactions with reactive metals; e.g. measure gas production; strong acids react at
higher rates than weaker acids.
8. 4. The pH scale
8.4.1 Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale.
8.4.2 Identify which of two or more aqueous solutions is more acidic or alkaline using pH values.
8.4.3 State that each change of one pH unit represents a 10-fold change in the hydrogen ion concentration
[H+(aq)].
The
pH scale changes
is used to
comparewhen
neutrality,
or basicity
of by
different
solutions.
8.4.4Deduce
in [H+(aq)]
the pH acidity
of a solution
changes
more than
one pH unit.
The pH is measured using a pH meter or pH paper; the latter consists of a mixture of indicators which
respond by showing different colours at different concentrations of H+ ; each indicator shows a
particular colour at a certain concentration of H+. The lower the pH the greater the acidity; the greater
the pH, the greater basicity and the lower the concentration of H+.
The pH can be derived from the concentration of the H+. As the [H+] is often a very small number the
negative logarithm is taken; a negative logarithm to end up with a positive number.
So:
pH = - log [H+]
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pOH = - log [OH-]
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Examples:
 an acidic solution with a hydrogen ion concentration of 1 x 10-3 has a pH of 3
 an alkali solution with a [H+] of 1 x 10-9 has a pH of 9.
So one pH unit represents a tenfold change in acidity or basicity !!!!
Example: a solution with a pH of 2 has:

a tenfold increase in [H+] than a solution with a pH of 3

100 times more [H+] than a solution with a pH of 4

but a tenfold decrease in [H+] than s solution with pH 1.
Use the following simulation:
 http://phet.colorado.edu/en/simulation/acid-base-solutions to reinforce ideas about strong and
weak acids and alkalis and pH.
 http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/acidbasepH/ph_
meter.html
 http://phet.colorado.edu/en/simulation/ph-scale
Other simulations: http://www.afn.org/~afn02809/#ACID-BASE%20CHEMISTRY%20RESOURCES
IB questions
1. (N07) Which acids are strong?
I. HCl (aq)
A. I and II only
II. HNO3 (aq)
B. I and III only
III. H2SO4(aq)
C. II and III only
D. I, II and III
2. (N07) The pH of a solution changes from pH = 1 to pH = 3. What happens to the [H+] during this
pH change?
A. It increases by a factor of 100.
B. It decreases by a factor of 100.
C. It increases by a factor of 1000.
D. It decreases by a factor of 1000
3. (N06) Which is a Brønsted-Lowry acid-base pair?
A. H2O and O2-
B. CH3COOH and CH3COO-
C. NH2- and NH4+
D. H2SO4 and SO42-
4. (N06) Which is not a strong acid?
A. Nitric acid
B. Sulfuric acid
C. Carbonic acid
D. Hydrochloric acid
5. (N06) Lime is added to a lake to neutralize the effects of acid rain. The pH value of the lake water
rises from 4 to 7. What is the change in concentration of H ions in the lake water?
A. An increase by a factor of 3
B. An increase by a factor of 1000
C. A decrease by a factor of 3
D. A decrease by a factor of 1000
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6 (M07) Solutions of hydrochloric acid (HCl (aq)) and ethanoic acid (CH3COOH (aq)) of the same
concentration reacted completely with 5.0 g of calcium carbonate in separate containers. Which
statement is correct?
A. CH3COOH(aq) reacted slower because it has a lower pH than HCl (aq).
B. A smaller volume of CO2 (g) was produced with CH3COOH(aq) than with HCl (aq).
C. A greater volume of CO2 (g) was produced with CH3COOH(aq) than with HCl (aq).
D. The same volume of CO2 (g) was produced with both CH3COOH(aq) and HCl (aq).
7. (N05) Lime was added to a sample of soil and the pH changed from 4 to 6. What was the
corresponding change in the hydrogen ion concentration?
A. increased by a factor of 2
B. increased by a factor of 100
C. decreased by a factor of 2
D. decreased by a factor of 100
8. (N05) When the following 1.0 mol dm-3 solutions are listed in increasing order of pH (lowest first),
what is the correct order?
A. HNO3 < H2CO3 < NH3 < Ba(OH)2
B. NH3 < Ba (OH)2 < H2CO3 < HNO3
C. Ba (OH)2 < H2 CO3 < NH3 < HNO3
D. HNO3 < H2CO3 < Ba (OH)2 < NH3
9. (M06) Which change in [H+] causes the biggest increase in pH?
10. (M06)
11. (M02) The ionisation of sulphuric acid is represented by the equations below:
H2SO4 (aq) + H2O (l)
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
H3O+ (aq)
+ HSO4- (aq)
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HSO4- (aq) + H2O (l) 
What is the conjugate base of HSO4-?
A. H2O (l)
B. H3O+ (aq)
H3O+ (aq)
+ SO4-2 (aq)
C. H2SO4 (aq)
D. SO4-2 (aq)
12. (M99) In which reaction below does the first species react as a Lewis acid?
13. (M99) In the equilibrium below
which represents a conjugate acid-base pair
14. (M00) In the equilibrium below, what are the two conjugate bases?
A. CH3COOH and H2O
C. CH3COOH and H3O+
B. CH3COO- and H3O+
D. CH3COO- and H2O
15. (N99) For the reaction below, a Bronsted-Lowry acid is
A. NH3 (aq) because it contains the largest number of hydrogen atoms.
B. NH3 (aq) because it accepts a proton from HNO2 (aq).
C. HNO2 (aq) because it has lone pairs of electrons on the oxygen atoms.
D. HNO2 (aq) because it donates a proton to NH3 (aq)
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16. Which statement about hydrochloric acid is false?
A. It can react with copper to give hydrogen.
B. It can react with sodium carbonate to give carbon dioxide
C. It can react with ammonia to give ammonium chloride
D. It can react with copper oxide to give water
17. 1.00 cm3 of a solution has a pH of 3. 100 cm3 of the same solution will have a pH of:
A. 1
B. 3
C. 5
D. impossible to calculate from the data given
18. Which statement(s) is/are true about separate solutions of a strong acid and a weak acid both with
the same concentration?
I. They both have the same pH
A. I and II
II. They both have the same electrical conductivity
B. I only
C. II only
D. Neither I nor II
19. Identify the correct statement about 25 cm3 of a solution of 0.1 mol dm3 ethanoic acid, CH3COOH.
A. It will contain more hydrogen ions than 25 cm3 of a solution of 0.1 mol dm3 hydrochloric acid
B. It will have a pH greater than 7
C. It will react exactly with 25 cm3 of a solution of 0.1 mol dm3 sodium hydroxide.
D. It is completely dissociated into ethanoate and hydrogen ions in solution.
20. (N02) When the following 0.10 mol dm3solutions are arranged in order of increasing pH (lowest
first), what is the correct order?
NH3 (aq),
HCl (aq),
NaOH (aq) ,
CH3COOH (aq).
A. NaOH , NH3 , CH3COOH , HCl
C. HCl , CH3COOH , NH3, NaOH
B. HCl , CH3COOH , NaOH, NH3
D. NaOH, NH3, HCl, CH3COOH
21. (M02) Solutions P, Q, R and S have the following properties:
P: pH = 8
Q : [H+] = 1 x 10-3 mol dm-3
R : pH = 5
S : [H+] = 2 x 10-7 mol dm-3
When these are arranged in order of increasing acidity (least acidic first), the correct order is
A. P,S,R,Q
B. Q,R,S,P,
C. S,R,P,Q,
D. R,P,Q,S
22. Which statements is true about 2 solutions one with pH of 3 and the other with a pH of 6?
A. The solution with a pH of 3 is twice as acidic as the solution with a pH of 6.
B. The solution with a pH of 6 is twice as acidic as the solution with a pH of 3.
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C. The hydrogen ion concentration in the solution with a pH of 6 is one thousand times greater
than that in the solution with a pH of 3.
D. The hydrogen ion concentration in the solution with a pH of 3 is one thousand times greater
than that in the solution with a pH of 6.
23. The difference between a weak and strong acid is that
A. solutions of weak acids cannot conduct an electric current but solutions of strong acids can
conduct an electric current.
B. concentrated solutions can be prepared from strong acids but not from weak acids.
C. the degree of dissociation is greater for strong acids than for weak acids.
D. weak acids are less soluble than strong acids.
24. (M00) 10 cm3 of an HCl solution with a pH value of 2 was mixed with 90 cm3 of water. What will be
the pH of the resulting solution?
A. 1
B. 3
C. 5
D. 7
25. (M89) Given the net ionic equation
H2O + PO43-  HPO42- + OHA. H2O is an acid and PO43- is its conjugate acid
B. H2O is an acid and OH- is its conjugate base
C. PO43- is a Bronsted acid because it is a proton acceptor
D. H2O is a Lewis acid because it is a proton donor.
PAPER 2
1. (N07)
(a) (i) A solution of hydrochloric acid has a concentration of 0.10 mol dm−3 and a pH value of 1. The
solution is diluted by a factor of 100. Determine the concentration of the acid and the pH
value in the diluted solution.
[2]
(ii) Explain why 0.10 mol dm-3 ethanoic acid solution and the diluted solution in (a) (i)
have similar [H+ ] values.
[3]
(b) Suggest one method, other than measuring pH, which could be used to distinguish between
solutions of a strong acid and a weak acid of the same concentration. State the expected
results.
[2]
2. (M07)
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3. (N05) The pH values of solutions of three organic acids of the same concentration were
measured.
(a) Identify which solution is the least acidic.
[1]
(b) Deduce how the [ H+] values compare in solutions of acids Y and Z.
[2]
(c) Arrange the solutions of the three acids in decreasing order of electrical conductivity,
starting with the greatest conductivity, giving a reason for your choice.
[2]
4. (N06) A titration was carried out to determine the concentration of 25.0 cm3 of an aqueous solution
of nitric acid. The pH value of the liquid in the flask was measured as 0.100 mol dm−3.
Aqueous sodium hydroxide was added. The results are shown on the graph below.
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(i) Use the graph to determine the value of [ H+] of the nitric acid solution.
[1]
(ii) Determine the pH value when the value of [H+] has decreased to 1 x 10-3 mol dm-3.
[1]
(iii) Use the graph to determine the volume of 0.100 mol dm-3 aqueous sodium hydroxide solution
needed to exactly neutralize the nitric acid.
[1]
(iv) Calculate the concentration, in mol dm-3, of the nitric acid.
[2]
(v) The pH values of three acidic solutions, X, Y and Z, are shown in the following table:
Solutions X and Z have the same acid concentration. Explain, by reference to both acids, why
they have different pH values.
[2]
(vi) Deduce by what factor the values of [H+] in solutions X and Y differ. [1]
5. (N02) Carbonic acid (H2CO3) is described as a weak acid and hydrochloric acid (HCl) is described
as a strong acid.
(a) Explain, with the help of equations, what is meant by strong and weak acid using the above
acids as examples.
[4]
(b) Outline two ways, other than using pH, in which you could distinguish between carbonic acid
and hydrochloric acid of the same concentration.
[2]
(c) A solution of hydrochloric acid, HCl (aq), has a pH of 1 and a solution of carbonic acid, H2CO3,
has a pH of 5. Determine the ratio of the hydrogen ion concentrations in these 2 solutions.
[2]
(d) The relative strengths of acids can be illustrated by the following equation.
(i) Identify the acid and its conjugate base and the base and its conjugate acid in the above
equation.
[2]
(ii) Name the theory that is illustrated in (d) (i).
[1]
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(e) Give examples of both a strong base and a weak base, clearly indicating which is which.
[2]
(f) Give an equation for the reaction between carbonic acid and one of the bases given in (e).
[3]
(g) Carbonic acid can be used to treat wasp (an insect) stings.
[3]
(i) Suggest what this indicates about the nature of wasp stings.
(ii) Name the type of reaction that occurs.
(iii) Explain why hydrochloric acid is not used to treat wasp stings.
6. (N01) Five unlabelled bottles are known to contain the following 0.10 aqueous solutions: 3 mol dm-3
NaCl,
NaOH,
HCl, CH3COOH, NH3
(a) Describe and explain how the pH values of these five solutions could be used to identify them. [5]
(b) Experiments were conducted to illustrate some properties of sodium hydrogen carbonate,
NaHCO3.
(i) In one experiment some solid was added to aqueous NaOH. After stirring NaHCO 3
the pH decreased to 9. Write a balanced chemical equation for the reaction and explain
the decrease in pH.
[2]
(ii) In another experiment solid was added to an aqueous solution of HCl. After NaHCO3
stirring the pH increased to 5. Write a balanced equation for the reaction and explain
this result.
(c) Describe how the two reactions of NaHCO3 in (b) illustrate the Brønsted–Lowry theory of acids
and bases.
[2]
(d) The graph below shows how the conductivity of a strong and a weak monoprotic acid change
as the concentration changes:
(i) Identify the strong acid and the weak acid from the above data. Give reasons for your
choices.
[6]
(ii) Describe how magnesium metal can be used to distinguish between solutions of the
strong acid and the weak acid of the same concentration.
[2]
(iii) Compare the volume of 0.10 mol dm-3 NaOH required to react exactly with 10.0 3 cm3
of 0.10mol dm-3 solutions of each of these acids.
[1
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7. (N00) 25.0 cm3 of hydrochloric acid of known concentration is titrated with a dilute sodium
hydroxide solution. The pH of the mixture is measured continuously as shown in the graph below:
(a) (i) From the graph, determine the pH after 10.0 of sodium hydroxide solution is cm 3
added.
[2]
(ii) Determine the concentration of hydrochloric acid before titration and state its units.
[1]
(iii) From the graph, determine the volume of sodium hydroxide solution required to
neutralise the hydrochloric acid.
[1]
(iv) Calculate the concentration of the sodium hydroxide solution and state its units.
[2]
(b) (i) Hydrochloric acid is a strong acid, whereas ethanoic acid is a weak acid. What is the
difference between a strong acid and a weak acid?
[2]
(ii) What mass of ethanoic acid would you use and how would you prepare 0.500 dm3 of a
0.500 mol dm3 ethanoic acid solution? (Mr of ethanoic acid = 60 0).
[2]
8. Using examples and balanced chemical equations (including states), define the terms weak base
and strong base. Discuss two different experimental methods to distinguish between 0.1 mol dm-3
solution each of a weak acid and a strong acid.
9. Explain the difference between strength and concentration. The pH of each 0.001 mol dm-3
hydrochloric acid and 0.056 mol dm-3 ethanoic acid is 3. Compare the concentrations of the two
acids, their strengths and explain the relationship between their hydrogen ion concentrations.
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Mark scheme
PAPER 1
1
2
3
4
5
6
7
8
D
B
B
C
D
D
D
A
9
10
11
12
13
14
15
16
D
D
D
B
D
D
D
A
17
18
19
20
21
22
23
24
B
D
C
C
A
D
C
B
25
26
27
28
29
30
31
32
B
PAPER 2
1. (N07)
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2. (M07)
3. (N05)
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4. (N06).
5. (N02)
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6. (N01)
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7. (N00)
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