C10 acids, bases and salts

LEARNING OUTCOMES
Define acid and acid anhydride
Investigate the reactions of non-oxidising acids with
metals, carbonates, hydrogen carbonates and bases
Define base and alkali
Investigate the reaction of bases with ammonium
salts
Relate acidity and alkalinity to the pH scale
Discuss the strength of acids and alkalis on the
basis of their completeness of ionisation
Define acidic, basic, amphoteric and neutral oxides
Chapter 10
Acids, Bases and Salts

LEARNING OUTCOMES
Define salt
Identify an appropriate method of salt preparation
based on the solubility of the salt
Distinguish between acidic and normal salts
Investigate neutralisation reactions using indicators
and temperature changes
Chapter 10

Chapter 10
What are acids?
 Fruits like apples, oranges and pineapples taste sour because they
contain acids.
 Acids also turn blue litmus paper red.
 Acids produce hydrogen ions H+
in water.

An acid is a substance which produces hydrogen ions, H+
(aq) in water.
Definition of An Acid
 For example, hydrochloric acid dissolves in
water to form hydrogen ions and chloride ions:
HCl(aq)  H+
(aq) + Cl-
(aq)
 It is the hydrogen ions which turn blue litmus
to red and give acids their characteristic properties.
Chapter 10

 Acids react with metals to produce
hydrogen gas.
E.g. Mg + H2SO4  MgSO4 + H2
Other chemical properties of acids
 Acids react with carbonates to
produce carbon dioxide.
E.g.
CaCO3 +2HCl  CaCl2 + H2O + CO2
Chapter 10
What are acids?
( test for hydrogen gas)
(test for carbon dioxide)
Limewater
turns chalky
HCl+CaCO3
pop

Other chemical properties of acids
 Acids react with bases to form a salt and water only.
 E.g. sulphuric acid reacts with copper(II) oxide to form a salt
called copper(II) sulphate and water:
H2SO4 + CuO  CuSO4 + H2O
 This reaction is called neutralisation.
What are acids?
Chapter 10

A Strong Acid
 A strong acid is an acid that
is completely ionised in
water. This means that all
the acid molecules become
ions in the water.
 Examples of strong acids are: sulphuric acid, hydrochloricExamples of strong acids are: sulphuric acid, hydrochloric
acid and nitric acid.acid and nitric acid.
Strong acid
Chapter 10

A Weak Acid
 E.g.s. of weak acids are: ethanoic acid, citric acid and
carbonic acid.
Weak acid
 A weak acid is an acid that is
only partially ionised in water.
This means that only a few
molecules of the acid become
ions in water.
Chapter 10

Some Common Acids
Name of acid Formula
Sulphuric acid H2SO4
Hydrochloric acid HCl
Nitric acid HNO3
Citric acid C6H8O7
Ethanoic acid (vinegar) CH3COOH
Chapter 10

 Ethanoic acid is used in vinegar for cooking and
to preserve food such as vegetables.
Uses of Acids
 Hydrochloric acid is used in the industry to remove
rust from metals before they are painted.
 Sulphuric acid is used to make fertilisers and
detergents.
 Citric acid is used in making fruit salts.
Chapter 10

Quick check 1
1. What ions do acids produce in water?
2. State three properties of acids.
3. Explain what is meant by a strong acid. Give one
example of a strong acid.
4. Explain what is meant by a weak acid. Give one example
of a weak acid.
5. Some dry citric acid crystals are placed on a dry piece of
litmus paper. Will there be a colour change? Explain your
answer.
Solution
Chapter 10

Solution to Quick check 1
1. Hydrogen ions
2. (a) Acids have a sour taste.
(b) Acids turn blue litmus to red.
(c) Acids react with metals to produce hydrogen.
3. A strong acid is an acid that is completely ionised in water. E.g.
sulphuric acid.
4. A weak acid is an acid that is only partially ionised in water. E.g.
ethanoic acid.
5. There will be no colour change because there is no water, so
the citric acid cannot form hydrogen ions.
Return
Chapter 10

13
Bases
 A base is an oxide or hydroxide of a metal.
 Examples of bases are:
sodium oxide, sodium hydroxide, copper(II)
oxide, copper(II) hydroxide, etc.
 A base reacts with an acid to form a salt and
water only.
E.g. CuO + H2SO4  CuSO4 + H2O
 This process is called neutralisation.
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14
 If a base is soluble in water, it is called an alkali.
Alkalis
 Sodium hydroxide is an alkali because it dissolves in
water to produce hydroxide ions:
NaOH(aq)  Na+
(aq) + OH−
(aq)
 An alkali is a soluble base which
produces hydroxide ions, OH−
(aq)
in water.
Chapter 10

15
 Copper(II) hydroxide is a base but not an alkali. This is
because it is insoluble in water and hence cannot produce
hydroxide ions in water.
Difference between base and alkali
BASE
ALKALICuO
MgO
Ca(OH)2
NaOH KOH NH3(aq)
Fe2O3
Cu(OH)2
 Is this true?
All alkalis are bases,
but not all bases are alkalis.
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16
 Alkalis have a bitter taste and soapy feel.
 Alkalis turns red litmus to blue.
Chemical properties of alkalis
 Alkalis react with acids to from salt and water
only.
E.g. 1. NaOH + HCl  NaCl + H2O
E.g. 2 2KOH + H2SO4  K2SO4 + 2H2O
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17
 Alkalis react with ammonium salts to produce ammonia gas.
 Ammonia gas is acidic, thus it turns red litmus paper blue.
 Ammonia gas is very soluble in water and gives out a pungent
smell.
E.g.1: NaOH + NH4Cl  NaCl + NH3 + H2O
Chemical properties of alkalis
Sodium hydroxide +
ammonium chloride
E.g. 2: Ca(OH)2 + 2NH4Cl  CaCl2 + 2NH3 + 2H2O
NH3 gas produced turns
red litmus blue
Chapter 10

18
 Sodium hydroxide and potassium hydroxide are used in
making soaps.
Uses of Bases
 Ammonia solution is used in window cleaners.
 Magnesium hydroxide is used in toothpastes to neutralise
the acid produced by bacteria.
 Calcium hydroxide (slaked lime) is used to neutralise
acids found in acidic soil.
Chapter 10

19
Some Common Alkalis
Name Chemical formula
Sodium hydroxide NaOH
Potassium hydroxide KOH
Calcium hydroxide Ca(OH)2
Ammonia solution
(ammonium hydroxide)
NH3(aq)
Chapter 10

20
Quick check 2
1. What is a base? Give 3 examples of bases.
2. Define what is an alkali. Give 3 examples of alkalis.
3. State 3 properties of alkalis.
4. Explain why iron(II) hydroxide is a base, but not an alkali.
5. Write balanced chemical equations for the following
reactions:
(a) potassium hydroxide + ammonium chloride
(b) calcium hydroxide + ammonium chloride
Solution
Chapter 10

21
1. A base is an oxide or hydroxide of a metal.
E.g. sodium oxide, copper(II) oxide, calcium hydroxide.
2. An alkali is a soluble base which produces hydroxide ions in water.
E.g. sodium hydroxide, potassium hydroxide, calcium hydroxide.
3. (i) Alkalis turn red litmus blue.
(ii) Alkalis react with acids to produce a salt and water.
(iii) Alkalis react with ammonium salts to produce ammonia.
4. Iron(II) hydroxide is a base, but not an alkali because it is insoluble
in water, so it cannot produce hydroxide ions in water.
5. (a) KOH + NH4Cl  KCl + H2O + NH3
(b) Ca(OH)2 + 2NH4Cl  CaCl2 + 2H2O + 2NH3
Return
Chapter 10

22
Indicators
 Indicators are substances which show different
colours in acidic and alkaline solutions.
 Litmus is a common indicator. It is red in acidic
solutions and blue in alkaline solutions.
 Other important indicators are shown in the table
on the next slide.
Chapter 10

23
Indicators
Indicator Colour in
strong
Acids
pH at which
colour
changes
Colour in
strong
alkalis
Methyl orange red pH 4 yellow
Litmus red pH 7 blue
Phenolphthalein colourless pH 9 pink
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24
 The pH of a solution tells us how acidic or alkaline a
solution is.
 The pH is a measurement of the hydrogen ion concentration
in a solution.
 The pH scale ranges from 0 to 14.
 The pH of a solution can be measured with a pH meter.
The pH Scale
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25
 The lower the pH, the more acidic the solution is.
 The higher the pH, the more alkaline the solution is.
 pH 7 is neutral.
 Distilled water, sugar solution and most salt solutions are
neutral (pH 7).
The pH Scale
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26
 The Universal Indicator consists of a mixture of dyes which
changes its colour in different pH solutions.
 We can use the Universal Indicator to tell us the
approximate pH of a solution.
 The Universal Indicator or pH paper changes its colour
according to the pH shown in the chart below.
The Universal Indicator
Box of pH paper with
colour chart
Chapter 10

27
Types of Oxides
 Elements burn or react with oxygen to form oxides.
 There are 4 types of oxides: acidic oxides, basic oxides, amphoteric
oxides and neutral oxides.
 An acidic oxide is an oxide of a non-metal. It dissolves in water to form an
acid. Acidic oxides react with alkalis to form salts .
 A basic oxide is an oxide of a metal. If soluble, it will dissolve in water to
form an alkali. Basic oxides react with acids to form salts.
 An amphoteric oxide is an oxide which can react with both acids and
alkalis to form salts.
 A neutral oxide does not react with either acids or alkalis.
Chapter 10

28
Types of Oxides
Chapter 10
Acidic Oxides Basic Oxides Amphoteric Oxides
CO2 , SO2
NO2 , NO
Na2O, CaO, K2O,
MgO, CuO
Al2O3 , PbO ,
ZnO
React with
alkalis to form
salts
React with acids to
form salts
React with both
acids & alkalis to
form salts
Neutral Oxides
H2O, CO ,
N2O
Do not react with
both acids &
alkalis
4 TYPES OF OXIDES

29
Quick check 3
1. Name 3 common indicators and their colour change in strong
acidic and strong alkaline solutions.
2. What is meant by the pH of a solution? What is the pH of :
(a) hydrochloric acid, (b) citric acid, (c) sodium chloride
solution, (d) sodium hydroxide solution?
3. What are the 4 types of oxides? Give one example of each
type of oxide.
4. What colours would you expect to see when the following
indicators are added to a solution of pH 5?
(a) litmus, (b) phenolphthalein, (c) methyl orange
Solution
Chapter 10

30
1. Litmus: red, blue;
Phenolphthalein: colourless, pink;
Universal Indicator: red, violet
2. The pH of a solution measures the acidity or alkalinity of a
solution. (a) 0 – 1, (b) 3 – 4, (c) 7, (d) 13 – 14.
3. Acidic oxides, basic oxides, amphoteric oxides and neutral
oxides. E.g. sulphur dioxide, sodium oxide, aluminium oxide,
water.
4. (a) litmus: red, (b) phenolphthalein: colourless,
(c) methyl orange: yellow
Return
Chapter 10

31
 A salt is formed when an acid is
neutralised by a base.
 A salt contains two parts:
 Metal part : cation (comes from the
base)
 Non-metal part : anion (comes from
the acid)
Salts
+Acid Base
Salt
Chapter 10

32
Examples of Salts
Base (alkali) Acid Salt formed
Sodium hydroxide Hydrochloric acid Sodium chloride
Potassium hydroxide Hydrochloric acid Potassium chloride
Sodium hydroxide Sulphuric acid Sodium sulphate
Potassium hydroxide Sulphuric acid Potassium sulphate
Calcium hydroxide Nitric acid Calcium nitrate
Ammonia solution Nitric acid Ammonium nitrate
Table 1
Chapter 10

33
 Sodium chloride is used as table salt and to preserve
meat and vegetables.
 Sodium chloride is electrolysed to obtain sodium and
chlorine in the industry.
 Ammonium nitrate and ammonium sulphate are used
as plant fertilisers.
Uses of Salts
 Magnesium sulphate, commonly called Epsom salt, is
used as a bath-salt.
Chapter 10

34
Methods of Preparing Salts
ACID + ALKALI  SALT + WATER
1. Action of acid on alkali
 This process is called neutralisation.
 To carry out the neutralisation of the acid
and alkali exactly, a method called titration
is used.
 The salts listed in Table 1 can be prepared
by the titration method.
Chapter 10

35
To prepare sodium nitrate by neutralisation (titration method)
Chapter 10
Sodium nitrate and water
(phenolphthalein as indicator)
burette
Pipette

36
Chapter 10
To prepare sodium nitrate by neutralisation (titration method)

37
ACID + BASE  SALT + WATER
2. Action of acid on insoluble base
 This method is used for bases which are insoluble in water.
 Examples of salts prepared by this method:
* copper(II) sulphate from copper(II) oxide and sulphuric acid:
CuO + H2SO4  CuSO4 + H2O
* zinc chloride from zinc oxide and hydrochloric acid:
ZnO + 2HCl  ZnCl2 + H2O
Chapter 10

38
Preparation of copper(II) sulphate (acid on insoluble base)
Chapter 10
Step 1 Place about 50 cm³ of dilute
sulphuric acid in a beaker and gently
warm the acid. Copper(II) oxide is added,
a little at a time, to the acid, until no more
can dissolve.
Equation: CuO + H2SO4  CuSO4 + H2O
Step 2 Filter off the excess copper(II) oxide
using a filter paper and funnel. Collect the
filtrate which contains copper(II) sulphate in
an evaporating dish.

39
Preparation of copper(II) sulphate (acid on insoluble base)
Chapter 10
Step 3 Evaporate the copper(II) sulphate solution until it is saturated.
Allow the hot solution to cool to form crystals.
Step 4 Filter off the copper(II) sulphate crystals formed and dry
them by pressing them between sheets of filter paper.

40
Eg.1 Sulphuric acid on sodium carbonate
H2SO4 + Na2CO3  Na2SO4 + H2O + CO2
Eg.2 Hydrochloric acid on calcium carbonate
2HCl + CaCO3  CaCl2 + H2O + CO2
 This method is similar to the previous method; instead of the
oxide, the carbonate is added in excess to the acid.
3. Action of acid on a carbonate
ACID + CARBONATE  SALT + WATER + CO2
Chapter 10

41
Eg.1 Sulphuric acid on zinc
H2SO4 + Zn  ZnSO4 + H2
Eg.2 Hydrochloric acid on magnesium
2HCl + Mg  MgCl2 + H2
 NOTE:
Only metals like magnesium, zinc and iron are suitable. Metals
like sodium, potassium and calcium are explosive with acids;
while metals like lead and copper are unreactive with acids.
4. Action of acid on a metal
ACID + METAL  SALT + HYDROGEN
Chapter 10

42
Making zinc sulphate (acid on metal)
Chapter 10
Can you describe how zinc sulphate is prepared with the aid of the diagrams?

43
5. Double Displacement (Precipitation method)
 This method is used to prepare insoluble salts.
 Two solutions are mixed together to produce a precipitate of
the insoluble salt which can then be filtered off from the mixture.
+
AD (s)
AB (aq) CD (aq)
CB (aq)
E.g. Lead(II) nitrate + Sodium chloride  Lead(II) chloride + Sodium nitrate
Pb(NO3)2(aq) + 2NaCl(aq)  PbCl2(s) + 2NaNO3(aq)
Chapter 10

44
 Silver chloride
AgNO3(aq) + HCl(aq)  AgCl(s) + HNO3(aq)
 Barium sulphate
Ba(NO3)2(aq) + H2SO4(aq)  BaSO4(s) + 2HNO3(aq)
 Copper(II) carbonate
CuSO4(aq) +Na2CO3(aq)  CuCO3(s) + Na2SO4(aq)
Other salts made by precipitation method
Chapter 10

45
Table of soluble and insoluble salts
Soluble salts Insoluble salts
All sodium, potassium and
ammonium salts
All carbonates except those of
sodium, potassium and ammonium
All nitrates None
All sulphates except those of calcium,
lead and barium
Calcium sulphate, lead(II) sulphate
and barium sulphate
All chlorides except those of silver
and lead
Silver chloride and lead(II) chloride
 This table will be useful to you when preparing salts
Chapter 10

46
Quick check 4
1. Define what is salt. Give an example of a soluble and insoluble
salt.
2. State 4 methods of making salts.
3. State whether the following salts are soluble or insoluble:
(a) sodium carbonate, (b) calcium chloride, (c) barium sulphate,
(d) lead(II) nitrate, (e) lead(II) chloride.
4. State the method you would choose to prepare the following
salts:
(a) potassium nitrate, (b) zinc nitrate, (c) magnesium sulphate,
(d) copper(II) carbonate.
For each method, state the chemicals you will need and
write a balanced chemical equation for the reaction.
Solution
Chapter 10

47
1. A salt is formed when an acid is neutralised by a base.
E.g. soluble salt: sodium chloride
E.g. insoluble salt: calcium sulphate
2. (a) Acid on metal, (b) acid on base, (c) acid on carbonate,
(d) precipitation method
3. Soluble: sodium carbonate, calcium chloride, lead(II) nitrate; Insoluble: lead(II)
chloride, barium sulphate
4. (a) potassium nitrate: titration method; potassium hydroxide and
nitric acid; KOH + HNO3  KNO3 + H2O
(b) zinc nitrate: acid on carbonate; nitric acid and zinc carbonate;
2HNO3 + ZnCO3  Zn(NO3)2 + H2O + CO2
(c) magnesium sulphate: acid on metal; magnesium and sulphuric acid;
Mg + H2SO4  MgSO4 + H2
(d) copper(II) carbonate: precipitation method;
copper(II) sulphate and sodium carbonate;
CuSO4(aq) + Na2CO3(aq)  CuCO3(s) + Na2SO4(aq)
Chapter 10
Return

48
 The state symbols in a chemical equation tell us about the state of each reactant
and product.
 The following are the state symbols used:
 Solid  (s)
 Liquid  (l)
 Gas  (g)
 Aqueous solution  (aq)
 Example: CaCO3(s) + 2HCl(aq)  CaCl2(aq) + H2O(l) + CO2(g)
 The above equation tells us that solid calcium carbonate reacts with a solution of
hydrochloric acid to produce liquid water and carbon dioxide gas.
State symbols in equations
Chapter 10

49
 Ionic equations are general equations which can apply to
any particular reaction.
 They represent ions taking part in a reaction, leaving out
those ions which do not react (spectator ions).
 They contain state symbols.
 Only solutions (aq) can form ions; gases, solids and liquids
do not ionise.
Writing ionic equations
Chapter 10

50
Steps in writing ionic equations
Step 3: Rewrite the equation with the final ions left:
H+
(aq) + OH-
(aq)  H2O(l)
EXAMPLE 1
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
Step 1: Break substances with (aq) into its ions:
H+
(aq) + Cl-
(aq) + Na+
(aq) + OH-
(aq)  Na+
(aq) + Cl-
(aq) + H2O (l)
Step 2: Remove similar ions from both sides of equation.
Chapter 10

51
EXAMPLE 2
2HCl(aq)+ CaCO3 (s)  CaCl2 (aq)+ H2O (l) + CO2 (g)
Step 1: Break those with (aq) into its ions:
2H+
(aq) + 2Cl-
(aq)+ CaCO3 (s)  Ca2+
(aq) + 2Cl-
(aq)+ H2O (l) + CO2 (g)
Step 2: Remove similar ions on both sides.
Step 3: Rewrite the equation with the ions left:
2H+
(aq) + CaCO3(s)  Ca2+
(aq)+ H2O(l) + CO2(g)
Chapter 10

52
EXAMPLE 3
Pb(NO3)2(aq)+ 2NaCl(aq)  PbCl2 (s)+ 2NaNO3 (aq)
Step 1: Break those with (aq) into its ions:
Pb2+
(aq) + 2NO3
-
(aq) + 2Na+
(aq) + 2Cl-
(aq)  PbCl2(s) + 2Na+
(aq) + 2NO3
-
(aq)
Step 2: Remove similar ions on both sides.
Step 3: Rewrite the equation with the ions left:
Pb2+
(aq) + 2Cl-
(aq) PbCl2(s)
Chapter 10

53
Quick check 5
Construct (i) a balanced chemical equation and (ii) an ionic
equation for each of the following reactions:
(1) Sulphuric acid + potassium hydroxide
(2) Nitric acid + sodium hydroxide
(3) Silver nitrate solution + sodium chloride solution
(4) Calcium carbonate + hydrochloric acid
(5) Magnesium + hydrochloric acid
Solution
Chapter 10

54
1. H2SO4(aq) + 2KOH(aq)  K2SO4(aq) + 2H2O(l)
H+
(aq) + OH-
(aq)  H2O(l)
2. HNO3(aq) + NaOH(aq)  NaNO3(aq) + H2O(l)
H+
(aq) + OH-
(aq)  H2O(l)
3. AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
Ag+
(aq) + Cl-
(aq)  AgCl(s)
4. CaCO3(s) + 2HCl(aq)  CaCl2(aq) + H2O(l) + CO2(g)
CaCO3(s) + 2H+
(aq)  Ca2+
(aq) + H2O(l) + CO2(g)
5. Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g)
Mg(s) + 2H+
(aq)  Mg2+
(aq) + H2(g)
Chapter 10
Return

55
1. http://www.sciencebyjones.com/acids_bases_salts.htm
2. http://ull.chemistry.uakron.edu/genobc/Chapter_09/
3. http://www.chem.ubc.ca/courseware/pH/index.html
To learn more about Acids, Bases and Salts,
click on the links below!
Chapter 10

C10 acids, bases and salts

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C10 acids, bases and salts