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KS4 Chemistry
Chemical Reactions
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Contents
Chemical Reactions
Introducing chemical reactions
Thermal decomposition
Displacement and precipitation
Neutralization
Redox
Summary activities
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What is a chemical reaction?
What is a chemical reaction?
A chemical reaction is a change that takes place when
one or more substances (called reactants) form one or
more new substances (called products).
chemical reaction
reactants
products
There are many different types of chemical reactions.
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Types of chemical reaction
How many types of chemical reaction can you name?
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Types of chemical reaction
oxidation &
reduction
neutralization
thermal
decomposition
chemical
reaction
exothermic &
endothermic
displacement:
metals
displacement:
non-metals
precipitation
reversible
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Exothermic and endothermic reactions
What are exothermic and endothermic reactions?
exothermic reactions give out energy – they get hot
 ex = out (as in ‘exit’)
 thermic = relating to heat
endothermic reactions take in energy – they get cold
 en = in (as in ‘entrance’)
Most chemical reactions are exothermic.
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Examples of exothermic reactions
Many exothermic reactions occur in the lab and in everyday
life. Can you think of six exothermic reactions?
Exothermic reactions
Burning wood on a fire
Burning petrol in a car
Burning gas on a gas hob
Reacting an acid and alkali together
Burning magnesium
Rotting compost
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Irreversible reactions
Most chemical reactions are considered to be irreversible
because the products cannot easily be changed back into
reactants.
For example, once magnesium has reacted
with hydrochloric acid, it is difficult to get the
magnesium back.
magnesium
+
hydrochloric
magnesium

acid
chloride
+ hydrogen
In equations for irreversible
reactions, reactants and
products are joined by a
‘one-way’ arrow.
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Reversible reactions
Although most chemical reactions are difficult to reverse,
there are some reactions that are fully reversible.
One of the best known reversible reactions occurs when
copper sulfate crystals are heated.
hydrated
copper sulfate
anhydrous
copper sulfate
+
water
CuSO4.5H2O
CuSO4
+
5H2O
In equations for reversible
reactions, reactants and
products are joined by a
‘two-way’ arrow.
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Equilibrium reactions
In some reversible reactions, the forward and backward
reactions largely occur in the same conditions and at the
same rate.
These reaction are said to be in equilibrium – there is no
overall change in the amount of products and reactants.
One of the most important equilibrium reactions occurs in
the production of ammonia in the Haber process:
nitrogen
+
hydrogen
ammonia
N2 (g)
+
3H2 (g)
2NH3 (g)
No matter how long the reaction is left, there will always be
a mixture of nitrogen, hydrogen and ammonia.
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Contents
Chemical Reactions
Introducing chemical reactions
Thermal decomposition
Displacement and precipitation
Neutralization
Redox
Summary activities
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Thermal decomposition
Thermal decomposition is a reaction in which a compound
is broken down by heat into two or more simpler substances.
Generally, the more reactive a metal, the harder it is to
decompose its compounds by heating.
For example:
silver carbonate:
decomposes on
gentle heating
calcium carbonate:
decomposes on
strong heating
increase in
reactivity of
metal
potassium carbonate: is not thermally
decomposed
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Thermal decomposition – easy or hard?
Compound
Decomposition
mercury oxide
easy
sodium oxide
hard
iron oxide
medium
silver oxide
easy
zinc oxide
medium
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potassium
sodium
calcium
magnesium
aluminium
zinc
iron
copper
mercury
silver
gold
increase in reactivity
How easy will these metal compounds be to decompose:
easy, medium or difficult?
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Thermal decomposition of carbonates
When metal carbonates are heated, they decompose to
produce metal oxides and carbon dioxide.
This reaction is performed industrially to make calcium oxide
(quicklime) from calcium carbonate (limestone):
calcium
carbonate

heat
calcium
oxide
+
carbon
dioxide
CaCO3

CaO
+
CO2
Quicklime is used to make concrete and calcium hydroxide
(slaked lime).
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Thermal decomposition of metal oxides
Most metal oxides are thermally stable – they do not
decompose when heated.
Oxides of the least reactive metals can be thermally
decomposed more easily. For example, mercury oxide
decomposes when heated strongly:
mercury
oxide
2HgO

heat

mercury condenses at
the top of the test tube,
where it is cooler
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mercury
+
oxygen
2Hg
+
O2
oxygen gas
escapes
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Contents
Chemical Reactions
Introducing chemical reactions
Thermal decomposition
Displacement and precipitation
Neutralization
Redox
Summary activities
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The reactivity series
potassium
sodium
calcium
magnesium
aluminium
zinc
iron
copper
mercury
silver
gold
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increase in reactivity
The reactivity series is a list of metals in order of their
reactivity.
The reactivity series can be
used to make predictions
about the reactivity of metals –
for example, how a metal will
react with oxygen, water and
acids.
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Displacement reactions: metals
A metal displacement reaction occurs when a metal is
added to a compound of a less reactive metal.
more
reactive
metal
+
compound
less
of less
 reactive
reactive
metal
metal
+
compound
of more
reactive
metal
The less reactive metal is displaced from the compound and
becomes elemental metal. The more reactive metal forms a
new compound.
A metal will always displace another
metal that is lower in the reactivity series.
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Displacement reactions – examples
Will magnesium react with copper chloride?
magnesium
+
copper
magnesium

chloride
chloride
+ copper
Magnesium is more reactive than copper, so it displaces
copper from its compound.
Will silver react with magnesium chloride?
silver
+
magnesium
chloride

no reaction
Silver is less reactive than magnesium, so it does not
displace magnesium from its compound.
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The Thermit process
The Thermit process is a
displacement reaction between
aluminium and iron (III) oxide.
magnesium
fuse
iron oxide
aluminium powder
(thermite)
aluminium
Al
+ iron oxide
aluminium

oxide
+ iron
+

+
Fe2O3
Al2O3
Fe
Aluminium is more reactive than iron and displaces it
from the oxide.
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The Thermit process
The reaction between aluminium and iron oxide is so
exothermic that the displaced iron melts.
The reaction is used to weld iron and steel together;
for example, railway tracks.
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Is there a displacement reaction?
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Displacement reactions: halogens
A halogen displacement reaction occurs when a halogen is
added to a metal halide containing a less reactive halogen.
The less reactive halogen is
displaced from the compound
and the more reactive halogen
bonds with the metal to form a
new metal halide.
F
Cl
decrease in
reactivity
Br
I
For example:
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fluorine
+
F2 (aq)
+
sodium
chloride

sodium
fluoride
2NaCl (aq)  2NaF (aq)
+
chlorine
+
Cl2 (aq)
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Displacement reactions of halogens
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Is there a displacement reaction?
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Precipitation reactions
When two aqueous solutions are mixed, they may react to
form a product that is insoluble in water. The solid is called a
precipitate and the reaction is called a precipitation
reaction.
To predict whether a precipitation reaction will occur,
information on the solubility of the products is needed.
What are the symbols for these physical states?
 solid
(g)
 liquid
(l)
 gas
(g)
 aqueous
(aq)
(dissolved in water)
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Precipitation reactions: sulfur
The precipitation reaction between solutions of sodium
thiosulfate and hydrochloric acid is often used to measure
rates of reaction.
sodium
sodium
hydrochloric
 chloride
thiosulfate +
acid
Na2S2O3
(aq)
+
2HCl
(aq)
Both reactants
are colourless.
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2NaCl

(aq)
+
sulfur
dioxide
+
SO2
(g)
+
sulfur
+
S
(s)
+ water
+
H2O
(l)
Sulfur is insoluble
and precipitates,
turning the
solution cloudy.
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Precipitation reactions: copper hydroxide
Many metal hydroxides are insoluble and can be formed by
precipitation reactions. For example:
copper (II)
sulfate
CuSO4
(aq)
+
+
copper (II)
ammonium

hydroxide
hydroxide
2NH4OH
(aq)
Copper (II) sulfate
solution is blue.
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
Cu(OH)2
(s)
+
ammonium
sulfate
+
(NH4)2SO4
(aq)
Copper (II) hydroxide
is insoluble and
forms a blue solid at
the bottom.
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Precipitation reactions: iron hydroxide
Iron (III) hydroxide is another insoluble metal hydroxide that
can be formed by a precipitation reaction.
iron (III)
chloride
+
FeCl3 (aq)
+
iron (III)
hydroxide
+
sodium
chloride
3NaOH (aq)  Fe(OH)3 (s)
+
3NaCl (aq)
sodium
hydroxide
Iron (III) chloride
solution is
yellow.
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
Iron (III) hydroxide is
insoluble and forms a
deep brown solid at
the bottom.
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Precipitation and solubility
To help work out whether a precipitate will form in a reaction,
there are some general rules about solubility.
Soluble
All compounds of sodium, potassium and ammonium.
All nitrates.
Most chlorides, except silver and lead chlorides.
Most sulfates, except lead, barium and calcium sulfates.
Insoluble
Most carbonates and hydroxides, except those
of sodium, potassium and ammonium.
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Predicting precipitation reactions
There are three steps to working out whether a precipitate
will be formed in a reaction:
Example
1. Write down the names
sodium chloride & lead nitrate
of the reactants.
2. Swap over the
non-metal.
sodium nitrate & lead chloride
3. Are the products
soluble or insoluble?
Lead chloride is insoluble
and will form a precipitate.
sodium
chloride
+
2NaCl (aq)
+
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lead
nitrate

sodium
nitrate
Pb(NO3)2 (aq)  2NaNO3 (aq)
+
lead
chloride
+
PbCl2 (s)
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Will there be a precipitation reaction? (1)
Will a precipitate be formed when sodium nitrate and
magnesium sulfate react?
1. Write down the names
of the reactants.
sodium nitrate &
magnesium sulfate
2. Swap over the
non-metal.
sodium sulfate &
magnesium nitrate
3. Are the products
soluble or insoluble?
Both products are soluble
so no precipitate will form.
sodium
sulfate
sodium
nitrate
+
magnesium

sulfate
2NaNO3 (aq)
+
MgSO4 (aq)  Na2SO3 (aq)
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+
magnesium
nitrate
+
Mg(NO3)2 (aq)
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Will there be a precipitation reaction? (2)
Will a precipitate be formed when sodium sulfate and
barium nitrate react?
1. Write down the names
of the reactants.
sodium sulfate &
barium nitrate
2. Swap over the
non-metal.
sodium nitrate &
barium sulfate
3. Are the products
soluble or insoluble?
Barium sulfate is insoluble
and will form a precipitate.
sodium
sulfate
+
Na2SO4 (aq)
+
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barium
nitrate

sodium
nitrate
Ba(NO3)2 (aq)  2NaNO3 (aq)
+
barium
sulfate
+
BaSO4 (s)
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Precipitation: true or false?
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Contents
Chemical Reactions
Introducing chemical reactions
Thermal decomposition
Displacement and precipitation
Neutralization
Redox
Summary activities
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Acids
What are acids? They are substances that:
 Have a pH below 7.
 Turn litmus red.
 Turn universal indicator yellow, orange or red.
 Form solutions containing H+ ions. The more H+ ions in
the solution, the stronger the acid.
Universal indicator pH scale
1 2 3 4 5 6
strong
acid
7 8 9 10 11 12 13 14
weak
acid
neutral
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Common acids
What are some common acids?
Acid
Formula
Strength
sulfuric acid
H2SO4
strong
hydrochloric acid
HCl
strong
nitric acid
HNO3
strong
ethanoic acid
(vinegar)
CH3COOH
weak
1 2 3 4 5 6
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Bases
What are bases? They are substances that:
 Have a pH above 7.
 Turn litmus blue.
 Turn universal indicator dark green, blue or purple.
 Are capable of neutralizing acids.
Universal indicator pH scale
1 2 3 4 5 6
7 8 9 10 11 12 13 14
weak
base
strong
base
neutral
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More about bases
Bases are usually oxides, hydroxides or carbonates of metals.
Ammonia is a base that doesn’t contain a metal.
Some bases are soluble in water – they are called alkalis.
bases
pH > 7
alkalis
soluble
bases
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Most carbonates and hydroxides are
insoluble in water, except those of
sodium, potassium and ammonium.
All alkaline solutions contain OH– ions.
The more OH– ions in the solution, the
stronger the alkali.
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Common bases
What are some common bases?
Base
Formula
Strength
sodium hydroxide *
NaOH
strong
potassium hydroxide *
KOH
strong
calcium hydroxide *
Ca(OH)2
strong
ammonia *
NH3
weak
calcium carbonate
CaCO3
weak
* = the base is also an alkali
1 2 3 4 5 6
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7 8 9 10 11 12 13 14
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Neutralization reactions
A neutralization reaction occurs when an acid reacts with
a base or alkali to produce a neutral solution of salt and
water.
acid
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+
alkali

salt +
water
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Naming salts
The salt formed in a neutralization reaction takes its name
from both the base and the acid.
The first part of the salt comes from the first part of the base,
for example:
 magnesium oxide

magnesium salt
 ammonium hydroxide 
ammonium salt
The second part of the salt comes from the acid:
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 sulfuric acid

sulfate
 hydrochloric acid

chloride
 nitric acid

nitrate
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What is the name of the salt?
What are the names of the salts formed from these bases
and acids?
Base
Acid
Salt
calcium hydroxide
hydrochloric acid
calcium chloride
magnesium oxide
nitric acid
magnesium nitrate
copper oxide
sulfuric acid
copper sulfate
aluminium hydroxide nitric acid
aluminium nitrate
potassium hydroxide sulfuric acid
potassium sulfate
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Neutralization reactions: hydroxides
When a hydroxide is mixed with an acid, OH– ions react
with H+ ions from the acid to form water:
H+
+
OH– 
H2O
For example:
potassium
hydroxide
+
KOH (aq)
+
calcium
hydroxide
Ca(OH)2 (aq)
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potassium
hydrochloric

chloride
acid
HCl (aq)
+
sulfuric
acid
+
water
H2O (aq)

KCl (aq)
+

calcium
sulfate
+
+ H2SO4 (aq) 
CaSO4(aq)
water
+ 2H2O (aq)
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Neutralization reactions: oxides
When an oxide is mixed with an acid, O2– ions react with
H+ ions from the acid to form water:
2H+
+

O2–
H2O
For example:
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calcium
oxide
+
CaO (aq)
+
copper
oxide
+
CuO (aq)
+
hydrochloric

acid
2HCl (aq)
sulfuric
acid
calcium
chloride
 CaCl2 (aq)

copper
sulfate
H2SO4 (aq)  CuSO4 (aq)
+
water
+
H2O (aq)
+
water
+
H2O (aq)
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Neutralization reactions: carbonates
When a carbonate is mixed with an acid, CO32– ions react
with H+ ions from the acid to form water and carbon dioxide:
2H+
+ CO32–  H2O + CO2
For example:
calcium
carbonate
CaCO3
(aq)
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+
calcium
nitric

nitrate
acid
+
water
+
+
2HNO3  Ca(NO3)2
(aq)
(aq)
+
H2O
(aq)
+
carbon
dioxide
CO2
(g)
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Complete the neutralization reaction
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Obtaining insoluble salts
How can insoluble salts be obtained?
1. The acid and alkali are mixed, and the salt forms by
precipitation.
2. The mixture of products is filtered, trapping the salt.
3. The salt is rinsed in distilled water, then left to dry.
For example, barium sulfate can be obtained by mixing
barium chloride with sulfuric acid.
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barium
hydroxide
+
BaOH2 (aq)
+
sulfuric
acid

barium
sulfate
H2SO4 (aq)  BaSO4 (s)
+
water
+
2H2O (aq)
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Obtaining soluble salts from bases
How can soluble salts be obtained following a reaction
between an acid and a base (insoluble)?
For example, obtaining copper sulfate from copper oxide and
sulfuric acid:
1. Copper oxide is added to sulfuric acid. When it is heated,
the copper oxide dissolves, forming a blue solution.
2. More copper oxide is added until no more will dissolve.
This means that all the acid has been used up.
3. The mixture is filtered to remove the excess copper oxide.
4. The filtrate is heated to evaporate some of the water.
When it cools, copper sulfate crystals will form.
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Obtaining soluble salts from alkalis (1)
How can soluble salts be obtained following a reaction
between an acid and an alkali (soluble)?
For example, the reaction between sodium hydroxide and
hydrochloric acid produces sodium chloride, which is soluble.
sodium
hydroxide
+
NaOH (aq)
+
hydrochloric

acid
sodium
chloride
+
water

NaCl (aq)
+
H2O (aq)
HCl (aq)
There is no obvious sign when the reaction is complete, so
an indicator must be used to show when the solution is
neutral.
This process is called titration.
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Obtaining soluble salts from alkalis (2)
To run the titration:
1. 25 cm3 of sodium hydroxide is added to a flask. Two drops
of the indicator phenolphthalein are added. This turns
pink.
2. Hydrochloric acid is added to the flask, a little at a time,
from a burette.
3. When all the alkali has reacted with the acid, the indicator
turns colourless. The amount of acid used is noted.
4. The experiment is repeated, but without adding the
indicator, as this makes the salt impure.
5. The salt solution is heated to evaporate the water.
Crystals of sodium chloride will remain.
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Matching reactants and salts
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Neutralization: true or false?
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Contents
Chemical Reactions
Introducing chemical reactions
Thermal decomposition
Displacement and precipitation
Neutralization
Redox
Summary activities
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What is redox?
What does redox mean?
reduction and oxidation
So far, these terms have been used to describe the gain and
loss of oxygen.
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Oxidation = adding
oxygen to a substance.
Reduction = removing
oxygen from a substance.
E.g. production of iron
oxide during rusting
E.g. extracting iron from iron
oxide in a blast furnace.
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Oxidation and ions
When a metal burns, it is oxidized to form a metal oxide.
What happens to the metal atoms during the reaction?
Metal atoms lose electrons to form positive ions. The oxygen
atoms accept these electrons and form negative ions.
For example, when magnesium burns to form magnesium
oxide, each magnesium atom loses 2 electrons and becomes
a magnesium ion with a +2 charge.
2 electrons
Mg
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O
Mg2+ O2-
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Oxidation and electron loss
What has happened to magnesium when it reacts with
oxygen?
2 electrons
O
Mg
Mg2+ O2-
 It has been oxidized.
 It has lost electrons by changing from Mg to Mg2+.
Oxidation is the loss of electrons.
Magnesium can lose electrons to substances other than
oxygen, e.g. when it reacts with chlorine or sulfur. These
reactions both involve Mg  Mg2+, so they are also oxidation.
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Reduction and ions
When iron is extracted from iron oxide (iron ore), the oxygen
is removed and the iron is said to be reduced.
When the oxygen is removed, 3 electrons are transferred
back to each iron ion, which become atoms.
O2O2O2-
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Fe3+
2 electrons
from each ion
Fe3+
O
Fe
O
O
Fe
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Reduction and electron gain
What has happened to iron when oxygen is removed?
O2O2O2-
Fe3+
2 electrons
from each ion
Fe3+
O
Fe
O
O
Fe
 It has been reduced.
 It has gained electrons by changing from Fe3+ to Fe.
Reduction is the gain of electrons.
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Electrons and OILRIG
An easy way to remember what happens to electrons during
oxidation and reduction is to think…
Oxidation
…OILRIG
Is
Loss …of electrons
Reduction
Is
Gain …of electrons
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Redox reactions
When a substance is oxidized, it loses electrons. Another
substance must gain these electrons and become reduced.
For example, when magnesium burns:
 magnesium loses electrons:
Mg  Mg2+ = oxidation
 oxygen gains electrons:
O  O2– = reduction
The overall reaction is reduction and oxidation = redox.
Oxidation and reduction
always take place together.
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Oxidized or reduced?
For each reaction, decide which product is oxidized and which
product is reduced.
calcium
+
oxidized
zinc oxide
reduced
aluminium
oxidized
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oxygen
 calcium oxide
reduced
+
hydrogen  zinc
+
water
oxidized
+ iron oxide  iron + aluminium oxide
reduced
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Oxidized or reduced?
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Contents
Chemical Reactions
Introducing chemical reactions
Thermal decomposition
Displacement and precipitation
Neutralization
Redox
Summary activities
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Glossary (part 1)
 displacement – A type of reaction in which a metal or
halogen replaces a less reactive metal or halogen in a
compound.
 equilibrium – A type of reaction in which products are
broken down at the same rate as reactants are combining.
 endothermic – A type of reaction that takes in energy.
 exothermic – A type of reaction that gives out energy.
 neutralization – A type of reaction in which an acid and
base react to form a salt and water.
 oxidation – A type of reaction involving the loss of
electrons.
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Glossary (part 2)
 precipitation – A type of reaction in which two aqueous
solutions react to form an insoluble product.
 reaction – A change that takes place when one or more
substances form one or more new substances.
 redox – A type of reaction in which oxidation and reduction
take place together.
 reduction – A type of reaction involving the gain of
electrons.
 decomposition – A type of reaction in which a substance
is broken down into two or more simpler substances.
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Anagrams
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Identify the reactions
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Multiple-choice quiz
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