Download Topic 8: ACIDS and BASES

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 that can be released easily as an ion, H+ (a
replaceable hydrogen). This happens when the hydrogen is part of a weak covalent bond or polar
covalent bond. If the acid has 1 replaceable hydrogen it is called a monoprotic acid.

The hydrogen ion 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 to form an hydroxonium ion.
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 . The following can act as Bronsted-Lowry bases:

substances such as oxides, hydroxides, carbonates and hydrogen carbonates that contain ions like
OH-, O2-, HCO32- or CO32- as these ions can accept the proton/H+ (s) from the acid (because they
contain atoms with 1 or more lone pairs) as shown below ;
IB chemistry topic 8
1|Page
OH- (aq) +
H+ (aq)
O2- (s) +

 H2O (l)
2H+ (aq)
CO32- (s) +
 H2O (l)
2H+ (aq)  H2O (l) + CO2 (g)
HCO3- (s) + H+ (aq)  H2O (l) + CO2 (g)
molecules that have lone pairs (e.g. ammonia and water); during an acid-base reaction, the Bronsted
–Lowry 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 + HCl  NH4+ + ClIn 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.
WHEN A BRONSTED LOWRY 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)
are reversible processes (which means an equilibrium exists).
This is the case even when it involves strong acid or strong bases like hydrochloric acid and sodium
hydroxide. In the case of strong acids and bases the reverse 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.
Therefore, in the equilibrium process there are 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:
IB chemistry topic 8
2|Page
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)

base
NH4+ (aq)
Cl- (aq)
+
conjugate acid
of the base NH3
conjugate base
of the acid HCl
Conjugate acid-base pairs from the above example:


HCl and ClNH3 and NH4+ .
A conjugate acid-base pair is a pair of species which differ by the presence of one H+ only; the acid
member of the pair always has an extra H+.
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 e.g. water, HSO 4- and HCO3-.
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- ?
B. H2SO3-
A. 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?
B. CH3
A. CH4
C. CH3+
D. CH3-
6. Which one can function as both an acid and a base according to the Bronsted-Lowry definition?
A. HS-
B. S2-
C. NH4+
D. Al3+
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-
IB chemistry topic 8
C. HPO42-
D. PO433|Page
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
 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;
IB chemistry topic 8
4|Page
 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 e.g. 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
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. OH-
B. N3-
D. NH4+
C. NH3
3. Classify each of the following species as a Lewis acid or base.
(a) CO2
(b) H2O
(c) I-
(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+
IB chemistry topic 8
(e) CN- + H2O  HCN + OH(f) SO2 + H2O  H2SO4
(g) HNO2 + OH-  NO2- + H2O
5|Page
(d) CH3COO- + 2HF  CH3COOH
+ 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 that 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)

Mg (s)
2NaHCO3 (s) +
H2SO4 (aq)
(net ionic equation for the first one =
 Na2SO4 (aq) +
2H2O (l) + CO2 (g)
CO32- (s) + 2H+ (aq)  CO2 (g) + 2H2O (l))
react with ammonia to form ammonium salts, e.g.:
NH3 (s) +

 Mg(NO3)2 (aq) + H2 (g)
react with metal carbonates and metal hydrogen carbonates to produce a salt, carbon dioxide and
water; examples:
 CaCO3 (s) + 2HCl (aq)
 CaCl2 (aq) + H2O (l) + CO2 (g)


as the sodium and sulfate ions are spectator ions.
react with the more reactive metals to produce a salt and hydrogen;
2HNO3 (aq) +

2H2O (l)
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.
IB chemistry topic 8
6|Page
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 symbol 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.
Strength of an acid
Strong acids ionize or dissociate (almost) completely (=equilibrium process which goes to virtual
completion) when they react with water.
In weak acids not all of the acid molecules ionize, in fact most acid molecules remain molecules; the
weaker the acid the fewer the number of acid molecules that ionize.
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.:
IB chemistry topic 8
HNO3

H+
+
NO3-
7|Page
!!!!! 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 ionizes readily or gives up very readily its hydrogen (either to water when it
dissolves or to a base).
Because of their different extent of ionization, 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.
NaOH (s) 
e.g.:
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)
2OH- (aq)
+
Ba2+ (aq)
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.
IB chemistry topic 8
8|Page
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)].
8.4.4Deduce changes in [H+(aq)] when the pH of a solution changes by more than one pH unit.
The pH scale is used to compare neutrality, acidity or basicity of different solutions.
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+]
pOH = - log [OH-]
Examples:
 an acidic solution with a hydrogen ion concentration of 1 x 10-3 mol dm-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 ([H+] = 1 x 10-2 mol dm-3) has:

a tenfold increase (= 10 times more) in [H+] than a solution with a pH of 3 ([H+] = 1 x 10-3 mol
dm-3).

100 times more [H+] than a solution with a pH of 4 ([H+] = 1 x 10-4 mol dm-3).

but a tenfold decrease (= 10 times less) in [H+] than a solution with a pH of 1 ([H+] = 1 x 10-1 mol
dm-3).
IB chemistry topic 8
9|Page
IB chemistry topic 8
10 | P a g e