2. 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
LEARNING OUTCOMES
3. 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
LEARNING OUTCOMES
4. Whatare 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.
5. Anacidisasubstancewhichproduceshydrogenions,H+(aq)in water.
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.
Definition ofAn Acid
6. Acids react with carbonates
to produce carbon dioxide.
E.g.
CaCO3 +2HCl CaCl2 + H2O +
CO2
Whatare acids?
Other chemical properties
of acids
Acids react with metals
to produce hydrogen gas.
E.g. Mg + H2SO4 MgSO4 +
H2
( testforhydrogen gas)
(testforcarbondioxide)
Limewate
r turns
chalky
HCl+CaCO3
pop
7. 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.
Whatare acids?
8. 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,
hydrochloric acid and nitric acid.
Strongacid
AStrongAcid
9. E.g.s. of weak acids are: ethanoic acid,
citric acid and carbonic acid.
Weakacid
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.
AWeakAcid
10. Name of acid
Sulphuric acid
Hydrochloric acid
Nitric acid
Citric acid
Ethanoic acid
(vinegar)
Formula
H2SO4
HCl
HNO3
C6H8O7
CH3CO
OH
SomeCommonAcids
11. 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.
Ethanoic acid is used in vinegar for cooking
and to preserve food such as vegetables.
Citric acid is used in making fruit salts.
Usesof Acids
12. 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
Quickcheck 1
13. 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
SolutiontoQuickcheck1
14. 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.
14
Bases
15. If abaseissolubleinwater,it iscalledanalkali.
Sodium hydroxide is an alkali because it
dissolves in water to produce hydroxide
ions:
NaOH(aq) Na+(aq) + OH−(aq) 15
An alkali is a soluble base
which produces hydroxide
ions, OH− (aq) in water.
Alkalis
16. 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.
BASE
16
CuO
Ca(OH)2
NaOH
ALKA
LI
KOH
NH3(aq)
Cu(OH)2
Fe2O3
MgO
Isthis true?
Allalkalisare bases,
butnotall basesare alkalis.
Differencebetweenbaseand alkali
17. Alkalis have a bitter taste and soapy feel.
Alkalis turns red litmus to blue.
Alkalis react with acids to from salt and water only.
E.g. 1. NaOH + HCl NaCl +
H2O
E.g. 2. 2KOH+ H2SO4 K2SO4 + 2H2O
17
Chemicalproperties of alkalis
18. E.g.1:
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.
Sodium hydroxide + ammonium chloride
E.g. 2: Ca(OH)2 + 2NH4Cl CaCl2
+ 2NH3 + 2H2O
NaOH + NH4Cl NaCl + NH3 +
H2O
NH3 gas produced turns red litmus blue
Chemicalproperties of alkalis
18
19. Sodium hydroxide and potassium hydroxide are
used in making soaps.
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.
19
Usesof Bases
21. 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
21
Quickcheck 2
22. 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
22
Solution to Quick check 2
23. 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.
Indicators
23
25. 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.
25
ThepH Scale
26. 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).
26
ThepH Scale
27. 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.
Box of pH paper
with colour
chart
TheUniversal Indicator
28. 27
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.
Typesof Oxides
29. 28
Typesof Oxides
BasicOxides
Acidic Oxides
CO2 , SO2
NO2 , NO
Na2O, CaO,
K2O, MgO,
CuO
AmphotericOxides
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
4TYPESOFOXIDES
30. 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
30
Quickcheck 3
31. 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
31
Solution to Quick check 3
32. 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)
+
Acid
32
Base
Salt
Salts
34. 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.
Magnesium sulphate, commonly called
Epsom salt, is used as a bath-salt.
34
Usesof Salts
35. ACID +ALKALI SALT + WA
TER
35
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.
Methodsof Preparing Salts
38. ACID +BASE SALT + WA
TER
38
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
Methodsof Preparing Salts
2. Actionofacidoninsoluble base
39. acid oninsoluble base
Preparation of copper(II) sulphate (acid on
insoluble base)
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
40. Preparationofcopper(II)sulphate(acidoninsoluble base)
acid oninsoluble base
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
41. Eg.1
41
Eg.2
Sulphuric acid on sodium carbonate
H2SO4 + Na2CO3
Na2SO4 + H2O + CO2
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.
ACID + CARBONA
TE SALT + WA
TER+CO2
Methodsof Preparing Salts
3. Actionofacidonacarbonate
42. Eg.1 Sulphuricacidon zinc
H2
SO4 + Zn ZnSO4 + H2
42
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.
ACID + METAL SALT + HYDROGEN
Methodsof Preparing Salts
4. Actionofacidona metal
47. 46
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
Quickcheck 4
48. 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) 48
SolutiontoQuickcheck4
Return
49. 49
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.
Statesymbolsin equations
50. 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.
50
Writingionic equations
51. 51
Writingionic equations
Steps in writing ionic equations 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. Step 3: Rewrite the equation
52. 52
Writingionic equations
Steps in writing ionic equations
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)
53. PbCl2(s)+2NaNO3(aq)
-
Step1:Breakthosewith(aq) intoits 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) 53
Writingionic equations
Stepsinwritingionic equations
EXAMPLE3
Pb(NO3)2(aq)+2NaCl(aq)
54. 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
54
Quickcheck 5
