Acids and bases
In this chapter learners will explore acid-base reactions and redox reactions. Redox
reactions were briefly introduced in gr10. The concepts of acids, bases, reduction, oxidation
and oxidation numbers are all introduced here. The following list provides a summary of the
topics covered in this chapter.

Acids and bases.
This chapter begins by revising all the concepts done on acids and bases up to this
point. Learners are reminded what an acid and a base are (in particular the BronstedLowry definition) and how the definition and concept have changed over time. Although
the most recent definition of an acid and a base is the Lowry definition this is not
covered at school level and the Bronsted-Lowry definition serves as a good working
model for the acids and bases that learners encounter at school.
The concept of a polyprotic acid is introduced although it is not in CAPS. This is done
to help learners understand how to handle acids such as sulfuric acid in reactions. You
should try to use polyprotic acids sparingly in your examples.

Conjugate acids and bases and amphoteric (amphiprotic) substances.
The concept of conjugate acids and bases requires learners to think about reactions
going in reverse. By writing the equation in reverse, learners can see how the acid
becomes a base. This base is said to be the conjugate base of the acid since it is
conjugated (linked) to the acid.

Acid-hydroxide, acid-oxide and acid-carbonate reactions.
Three different types of bases are examined in detail to see how they react with acids.
Several examples of each type are given and the general equation for the reaction is
also given.

Oxidation numbers for compounds.
This topic is placed after redox reactions in CAPS but must be taught before redox
reactions and so is placed before redox reactions in this book. This topic provides the
tools needed to understand redox reactions.

Balancing redox reactions.
In grade 10 learners learnt how to balance chemical equations by inspection. In this
topic they will learn how to balance redox reactions which often cannot be balanced by
inspection. The simpler examples can be balanced by inspection and this can be used
as a comparison for the two techniques. Learners need to be able to break a reaction
up into two parts and follow different chemical species through an equation. This skill
starts with conjugate acids and bases and carries over into this topic.
Coloured text has been used as a tool to highlight different parts of reactions. Ensure that
learners understand that the coloured text does not mean there is anything special about
that part of the reaction, this is simply a teaching tool to help them identify the important
parts of the reaction.
It is also important to note that this chapter is split across term 3 and term 4. Acids and
bases should be completed in term 3 and redox reactions are done in term 4.
All around you there are chemical reactions taking place. Green plants are
photosynthesising, car engines are relying on the reaction between petrol and air and your
body is performing many complex reactions. In this chapter we will look at two common
types of reactions that can occur in the world around you and in the chemistry laboratory.
These two types of reactions are acid-base reactions and redox reactions.
Acids and bases
What are acids and bases?
Activity 1: Household acids and bases
Look around your home and school and find examples of acids and bases. Remember that
foods can also be acidic or basic.
Make a list of all the items you find. Why do you think they are acids or bases?
Some common acids and bases, and their chemical formulae, are shown in Table Table 1.
Acid
Formula
Base
HCl
Hydrochloric acid
Formula
NaOH
Sodium hydroxide
H2SO4
Sulfuric acid
KOH
Potassium hydroxide
H2SO3
Sulfurous acid
Na2CO3
Sodium carbonate
CH3COOH
Acetic (ethanoic) acid
Ca(OH)2
Calcium hydroxide
H2CO3
Carbonic acid
Mg(OH)2
Magnesium hydroxide
HNO3
Nitric acid
NH3
Ammonia
H3PO4
Phosphoric acid
NaHCO3
Sodium bicarbonate
Table 1: Some common acids and bases and their chemical formulae.
Most acids share certain characteristics, and most bases also share similar characteristics.
It is important to be able to have a definition for acids and bases so that they can be
correctly identified in reactions.
Defining acids and bases
One of the first things that was noted about acids is that they have a sour taste. Bases were
noted to have a soapy feel and a bitter taste. However you cannot go around tasting and
feeling unknown substances since they may be harmful. Also when chemists started to
write down chemical reactions more practical definitions were needed.
A number of definitions for acids and bases have developed over the years. One of the
earliest was the Arrheniusdefinition. Arrhenius (1887) noticed that water dissociates (splits
up) into hydronium (H3O+) and hydroxide (OH−) ions according to the following equation:
2H2O (l)→H3O+(aq)+OH−(aq)
Tip:
For more information on dissociation, refer to Grade 10 (chapter 18: reactions in aqueous
solution).
Arrhenius described an acid as a compound that increases the concentration of H3O+ ions
in solution and abase as a compound that increases the concentration of OH− ions in
solution.
Look at the following examples showing the dissociation of hydrochloric acid and sodium
hydroxide (a base) respectively:
1. HCl (aq)+H2O(l)→H3O+(aq)+Cl−(aq)
Hydrochloric acid in water increases the concentration of H3O+ ions and is therefore
an acid.
2. NaOH (s)⟶H2ONa+(aq)+OH−(aq)
Sodium hydroxide in water increases the concentration of OH− ions and is therefore
a base.
Note that we write ⟶H2O to indicate that water is needed for the dissociation.
However, this definition could only be used for acids and bases in water. Since there are
many reactions which do not occur in water it was important to come up with a much
broader definition for acids and bases.
In 1923, Lowry and Bronsted took the work of Arrhenius further to develop a broader
definition for acids and bases. The Bronsted-Lowry model defines acids and bases in
terms of their ability to donate or accept protons.
Definition 1: Acids
A Bronsted-Lowry acid is a substance that gives away protons (hydrogen
cations H+), and is therefore called a proton donor.
Definition 2: Bases
A Bronsted-Lowry base is a substance that takes up protons (hydrogen
cations H+), and is therefore called a proton acceptor.
Below are some examples:
1. HCl (aq)+NH3(aq)→NH+4(aq)+Cl−(aq)
We highlight the chlorine and the nitrogen so that we can follow what happens to these
two elements as they react. We do not highlight the hydrogen atoms as we are
interested in how these change. This colour coding is simply to help you identify the
parts of the reaction and does not represent any specific property of these elements.
H\textcolor{red}{Cl}
(aq)+\textcolor{blue}{N}H3(aq)→\textcolor{blue}{N}H+4(aq)+\textcolor{red}{
Cl}−(aq)
In order to decide which substance is a proton donor and which is a proton acceptor,
we need to look at what happens to each reactant. The reaction can be broken down
as follows:
H\textcolor{red}{Cl} (aq)→\textcolor{red}{Cl}−(aq) and
\textcolor{blue}{N}H3(aq)→\textcolor{blue}{N}H+4(aq)
From these reactions, it is clear that HCl is a proton donor and is therefore an acid,
and that NH3 is aproton acceptor and is therefore a base.
2. CH3COOH (aq)+H2O (l)→H3O+(aq)+CH3COO−(aq)
Again we highlight the parts of the reactants that we want to follow in this reaction:
\textcolorredCH3COOH (aq)+H2\textcolor{blue}{O}
(l)→H3\textcolor{blue}{O}+(aq)+\textcolorredCH3COO−(aq)
The reaction can be broken down as follows:
\textcolorredCH3COOH (aq)→\textcolorredCH3COO−(aq) and
H2\textcolor{blue}{O} (l)→H3\textcolor{blue}{O}+(aq)
In this reaction, CH3COOH (acetic acid or vinegar) is a proton donor and is therefore
the acid. In this case, water acts as a base because it accepts a proton to form H3O+.
3. NH3(aq)+H2O (l)→NH+4(aq)+OH−(aq)
Again we highlight the parts of the reactants that we want to follow in this reaction:
\textcolor{blue}{N}H3(aq)+H2\textcolor{red}{O}
(l)→\textcolor{blue}{N}H+4(aq)+\textcolor{red}{O}H−(aq)
The reaction can be broken down as follows:
H2\textcolor{red}{O} (l)→\textcolor{red}{O}H−(aq) and
\textcolor{blue}{N}H3(aq)→\textcolor{blue}{N}H+4(aq)
Water donates a proton and is therefore an acid in this reaction. Ammonia accepts the
proton and is therefore the base.
Notice in these examples how we looked at the common elements to break the reaction into
two parts. So in the first example we followed what happened to chlorine to see if it was part
of the acid or the base. And we also followed nitrogen to see if it was part of the acid or the
base. You should also notice how in the reaction for the acid there is one less hydrogen on
the right hand side and in the reaction for the base there is an extra hydrogen on the right
hand side.
Amphoteric substances
In examples 2 and 3 above we notice an interesting thing about water. In example 2 we find
that water acts as a base (it accepts a proton). In example 3 however we see that water
acts as an acid (it donates a proton)!
Depending on what water is reacting with it can either react as a base or as an acid. Water
is said to beamphoteric. Water is not unique in this respect, several other substances are
also amphoteric.
Definition 3: Amphoteric
An amphoteric substance is one that can react as either an acid or base.
When we look just at Bronsted-Lowry acids and bases we can also talk about amphiprotic
substances which are a special type of amphoteric substances.
Definition 4: Amphiprotic
An amphiprotic substance is one that can react as either a proton donor (BronstedLowry acid) or as a proton acceptor (Bronsted-Lowry base). Examples of
amphiprotic substances include water, hydrogen carbonate ion (HCO−3) and
hydrogen sulfate ion (HSO−4).
Note: You may also see the term ampholyte used to mean a substance that can act as
both an acid and a base. This term is no longer in general use in chemistry.
Polyprotic acids [NOT IN CAPS]
A polyprotic (many protons) acid is an acid that has more than one proton that it can
donate. For example sulfuric acid can donate one proton to form the hydrogen sulfate ion:
Or it can donate two protons to form the sulfate ion:
In this chapter we will mostly consider monoprotic acids (acids with only one proton to
donate). If you do see a polyprotic acid in a reaction then write the resulting reaction
equation with the acid donating all its protons.
Some examples of polyprotic acids are: H2SO4, H2SO3, H2CO3 and H3PO4.
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Exercise 1: Acids and bases
Problem 1:
Identify the Bronsted-Lowry acid and the Bronsted-Lowry base in the following reactions:
1. HNO3(aq)+NH3(aq)→NO−3 (aq)+NH+4 (aq)
2. HBr (aq)+KOH (aq)→KBr (aq)+H2O (l)
Practise more questions like this
Answer 1:
1. We break the reaction into two parts:
HNO3 (aq)→NO−3(aq) and
NH3(aq)→NH+4(aq)
From this we see that the Bronsted-Lowry acid is HNO3 and the Bronsted-Lowry base
is NH3.
2. We break the reaction into two parts:
HBr (aq)→KBr (aq) and
KOH (aq)→H2O (l)
From this we see that the Bronsted-Lowry acid is HBr and the Bronsted-Lowry base
is KOH.
Problem 2:
1. Write a reaction equation to show HCO−3 acting as an acid.
2. Write a reaction equation to show HCO−3 acting as an base.
3. Compounds such as HCO−3 are …
Practise more questions like this
Answer 2:
1. HCO−3 (aq)→CO2−3(aq)+H+(aq)
2. HCO−3(aq)+H+(aq)→H2CO3(aq)
3. Amphoteric
Conjugate acid-base pairs
Look at the reaction between hydrochloric acid and ammonia to form ammonium and
chloride ions (again we have highlighted the different parts of the equation):
\textcolor{red}{HCl}
(aq)+\textcolorblueNH3(aq)→\textcolorblueNH+4(aq)+\textcolorredCl−(aq)
We look at what happens to each of the reactants in the reaction:
HCl (aq)→Cl−(aq) and
NH3(aq)→NH+4(aq)
We see that HCl acts as the acid and NH3 acts as the base.
But what if we actually had the following reaction:
\textcolorblueNH+4(aq)+\textcolorredCl−(aq)→\textcolor{red}{HCl}
(aq)+\textcolorblueNH3(aq)
This is the same reaction as the first one, but the products are now the reactants.
Now if we look at the what happens to each of the reactants we see the following:
NH+4(aq)→NH3(aq) and
Cl−(aq)→HCl (aq)
We see that NH+4 acts as the acid and Cl− acts as the base.
Tip:
Up to now you have looked at reactions as starting with the reactants and going to the
products. For acids and bases we also need to consider what happens if we swop the
reactants and the products around. This will help you understand conjugate acid-base pairs.
When HCl (the acid) loses a proton it forms Cl− (the base). And that when Cl− (the base)
gains a proton it forms HCl (the acid). We call these two species a conjugate acid-base
pair. Similarly NH3 and NH+4 form a conjugate acid-base pair.
Tip:
The word conjugate means coupled or connected.
We can represent this as:
Activity 2: Conjugate acid-base pairs
Using the common acids and bases in Table Table 1, pick an acid and a base from the list.
Write a chemical equation for the reaction of these two compounds.
Now identify the conjugate acid-base pairs in your chosen reaction. Compare your results to
those of your classmates.
Exercise 2: Acids and bases
Problem 1:
In each of the following reactions, label the conjugate acid-base pairs.
1. H2SO4(aq)+H2O (l)→H3O+(aq)+HSO−4(aq)
2. NH+4(aq)+F−(aq)→HF(aq)+NH3(aq)
3. H2O (l)+CH3COO−(aq)→CH3COOH (aq)+OH−(aq)
4. H2SO4(aq)+Cl−(aq)→HCl (aq)+HSO−4(aq)
Practise more questions like this
Answer 1:
1.
2.
3.
4.
Problem 2:
Given the following reaction:
H2O (l)+NH3(aq)→NH+4(aq)+OH−(aq)
1. Write down which reactant is the base and which is the acid.
2. Label the conjugate acid-base pairs.
3. In your own words explain what is meant by the term conjugate acid-base pair.
Practise more questions like this
Answer 2:
1. We break the reaction into two parts:
H2O (aq)→OH−(aq) and
NH3(aq)→NH+4(aq)
From this we see that the Bronsted-Lowry acid is H2O and the Bronsted-Lowry base
is NH3.
2.
3. A conjugate acid-base pair is a reactant and product pair that is transformed into each
other through the loss or gain of a proton. So for example an acid loses a proton to
form a base. The acid and the resulting base are said to be a conjugate acid-base pair.