Download KS4-Chemical-Reactions

Chemical Reactions
© Boardworks Ltd 2001
Oxidation
and reduction
Thermal
decomposition
Types of
chemical change
neutralisation
precipitation
Reversible
reactions
Exothermic
and endothermic
Displacement
Reactions:
• metals
Displacement
Reactions:
• non-metals
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Thermal Decomposition
Potassium carbonate is not thermally decomposed.
Calcium carbonate decomposes on strong heating
Silver carbonate decomposes on gentle heating
Gets harder
• A thermal decomposition is when heat causes a
chemical to break down to simpler substances.
• Compounds – but not elements - undergo thermal
decomposition.
• For compounds that contain metals we usually find:
the more reactive the metal, the harder it is to
decompose its compounds. Eg.
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• Generally, the more reactive the metal
is, the more difficult it is to decompose
it’s compounds.
• Fill in the last column: easy, medium or
hard
Compound
Mercury oxide
Sodium oxide
Iron oxide
Silver oxide
Zinc oxide
How easy to
decompose
easy
hard
medium
Potassium
sodium
calcium
magnesium
aluminium
zinc
iron
copper
mercury
silver
gold
Increasing reactivity
Thermal Decomposition
easy
medium
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Thermal Decomposition of carbonates
limestone
• When carbonates are
heated they release
carbon dioxide.
• This reaction is performed
industrially to make
calcium oxide (quicklime)
from calcium carbonate
(limestone). Quicklime is
used to make concrete
and to make calcium
hydroxide (slaked lime).
Calcium
Carbonate

Calcium
oxide
+
Waste
air and
carbon
dioxide
Hot air
1500°C
Carbon
dioxide
Calcium oxide (lime)
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Thermal Decomposition of metal oxides
• Most metal oxides are
thermally stable. (I.e. Do
not decompose when
heated.)
• Oxides of the least
reactive metals are
thermally decomposed.
• E.g. silver oxide begins to
break up at about 160oC
and mercury oxide
decomposes when
heated strongly.
Mercury
Oxide

Mercury
+
Mercury metal
and oxygen
formed
O O
O
O
Hg
Hg
Hg
Hg
Heat
Hg O Hg O
O Hg O Hg
Hg O Hg O
O Hg O Hg
Hg O Hg O
O Hg O Hg
Hg O Hg O
O Hg O Hg
Mercury oxide
decomposes
oxygen
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Exothermic and Endothermic Reactions
•Ex = out (as in exit.)
•En = in (as in entrance.)
• Exothermic reactions give out heat (get
hot.)
• Endothermic reactions take in heat (get
cold.)
• Many chemical reactions need some
energy to get them started (activation
energy) but then the majority of
chemical reactions are exothermic.
Shuttle fuel
burning.
Highly
exothermic
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Exothermic and Endothermic Reactions
• It is hard to think of examples of endothermic
reactions but there are lots of exothermic ones that
occur in the laboratory and in everyday life.
• List 8 exothermic reactions.
Some examples of exothermic reactions
Burning wood on a fire
Burning petrol in a car
Burning butane in a cigarette lighter
Burning gas in a gas hob
Reacting an acid and alkali together
Burning magnesium
Rotting compost etc etc
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• These are reactions where two
metals are competing to be
combined with a non-metal.
• The more reactive metal wins the
competition and becomes part of
a compound.
• The less reactive metal is
displaced and so is present as
the metal at the end of the
reaction.
Potassium
sodium
calcium
magnesium
aluminium
zinc
iron
copper
silver
gold
Increasing reactivity
Displacement Reactions: Metals
A more reactive metal (higher in the reactivity series) will
displace a less reactive metal from its compound
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Displacement Reactions: Metals
• Copper is quite low in the activity series.
• Several metals will displace it from its compounds.
magnesium
K
Na
Ca
Mg
Al
Zn
Fe
Cu
Ag
Au
Magnesium
sulphate solution
copper
sulphate
solution
Copper metal
Magnesium
more
reactive
+
Copper
sulphate
less
reactive

Magnesium
sulphate
+
Copper
Magnesium wins the competition.
Copper is displaced
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Displacement Reactions: Metals
• Here are some actual photos.
• The colour changes from blue to red/black as
copper metal is displaced.
K
Na
Ca
Mg
Al
Zn
Fe
Cu
Ag
Au
Photograph
at start of
reaction
Magnesium +
more
reactive
Photograph
at end of
reaction
Copper 
sulphate
less
reactive
Magnesium
sulphate
+
Copper
Magnesium wins the
competition. Copper is displaced
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Displacement Reactions: Metals
• The thermit reaction takes
place between aluminium
and iron oxide. It is so
exothermic that molten
iron is produced and the
reaction is used to repair
K
Na broken railway tracks.
Ca
Mg
Al
Zn
Fe
Cu
Ag
Au
Aluminium
more
reactive
+
Iron
Oxide
less
reactive

Aluminium
Oxide
Magnesium
fuse
iron oxide +
aluminium
powder
+
Iron
Aluminium wins the competition.
Iron is displaced and melts at the
high temperatures produced.
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Displacement Reactions: Metals
• Here is a photo of the thermit
reaction being carried out in a
laboratory
Magnesium
fuse
iron oxide +
aluminium
powder
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Activity
• Predict which mixtures will result in a reaction
Metal 
Solution
Iron
Iron
chloride
Magnesium
nitrate
Zinc nitrate
Copper
sulphate
Magnesium
Yes
No
No
Yes
Yes
Yes
Zinc
Copper
Yes
No
No
No
No
Yes
© Boardworks Ltd 2001
Displacement Reactions: Halogens
•
Fluorine
Chlorine
Bromine
Iodine
Increasing reactivity
• These are displacement
reactions where two
halogens are competing to
be combined with a metal.
• It is the more reactive
halogen that will win and
become part of a
compound.
• The less reactive halogen
remains (or becomes) the
element.
We can often tell which halogen is present from the
colour of the solution.
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Displacement Reactions: Halogens
• Eg. If chlorine solution is added to sodium
bromide
chlorine
solution
F
Cl
Br
I
At
Sodium chloride
solution
Sodium
bromide
solution
Bromine
Chlorine
more
reactive
+
Sodium
Bromide
less
reactive

Sodium
Chloride
+
Bromine
Chlorine wins the competition.
Bromine (red) is displaced
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Activity
• Predict what colour these will be after mixing.
• (The compounds of the halogens with Group 1
metals are all colorless.)
Halogen
Halide 
Chlorine
solution
Potassium
chloride
Bromine
solution
Br2
Potassium
bromide
Br2
Potassium
Iodide
I2
Iodine
Solution
I2
I2
I2
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Displacement Reactions: Halogens
• When writing equations for halogen
displacement reactions you must
remember that – when in the form of the
element – halogens exist in pairs.
• Eg. For chlorine and sodium bromide:
Chlorine +
Sodium
sodium

bromide
chloride
F
Cl
Br
I
At
+ bromine
Cl2(aq) + 2NaBr(aq)  2NaCl(aq) + Br2(aq)
Cl More
reactive
Br Less
reactive
Solution goes
yellow/brown as
bromine is produced.
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Activity
• Predict whether or not a chemical reaction
will occur,
• If no reaction - not write “no reaction.”
• Where there is a reaction write the names of
the products and then write a chemical
equation underneath..
1) iodine + sodium bromide solution 
No reaction
2) bromine + sodium chloride solution 
No reaction
3) chlorine + sodium iodide solution 
Sodium
chloride
Cl2(g) + 2NaI(aq)

F
Cl
Br
I
At
+ iodine
2NaCl(aq) + I2(aq)
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Reversible and Irreversible Reaction
• Most Chemical reactions are considered
irreversible in that the new products are not
readily changed back into reactants. Eg. Once
you have reacted magnesium with hydrochloric
acid it is very hard to get the magnesium back.
•
In the equations for irreversible reactions
reactants and products are joined by a “oneway arrow.”
magnesium + hydrochloric  magnesium + hydrogen
acid
chloride
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Reversible Reactions
• Although most chemical reactions are difficult to
reverse it is possible to find reactions ranging
from irreversible through to the fully reversible.
• One of the best known reversible processes is
heating copper sulphate. Note the double arrow
symbol in the chemical equation
Heat
Hydrated copper
sulphate
CuSO4.5H20
These decompose
anhydrous copper
sulphate
CuSO4
+
steam
5H2O
These combine
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Equilibrium Reactions.
• There are some reactions in which both the
“forward and backward” reactions occur to a
substantial extent under the same conditions.
• These lead to equilibrium mixtures of reactants
and products.
• One of the most important of these reactions
occurs in the Haber Process.
N2(g) + 3H2(g)
2 NH3(g)
However long you leave the reaction going you still get a
mixture of nitrogen, hydrogen and ammonia.
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Getting more product at equilibrium.
• There are some simple rules that can be used to move
the position of an equilibrium towards reactants or
products:
1. Exothermic reactions give more product at lower
temperatures. (Endothermic – the opposite)
2. Increasing the pressure in gas reactions favours
whichever side of the chemical equation has least gas
molecules.
What conditions will favour formation of more ammonia?
3H2(g) + N2 (g)
Low temperature
 2NH3 (g) (exothermic)
High pressure
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Precipitation Reactions
• A precipitation reaction is any reaction that produces an
insoluble compound when two aqueous solutions are
mixed.
• It is impossible to predict whether or not we will get
precipitation reactions unless we know something about
the physical states (especially solubility) of the various
reactants and products.
•
Here are the symbols that we use in chemical
equations to say what the physical state is:
–(s)
–(l)
–(g)
–(aq)
solid
liquid
gas
aqueous (dissolved in water)
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Precipitation Reactions – 1st example
• A precipitation reaction that is often used to
measure reaction rates occurs between sodium
thiosulphate and hydrochloric acid.
Sodium +
thiosulphate
hydrochloric
acid
aqueous
aqueous
Both reactants are
colourless and
dissolved (aq)
 sodium + sulphur + water + sulphur
chloride
dioxide
solid
aqueous solid
liquid
gas
Sulphur is
insoluble and
precipitates. This
makes the solution
go cloudy.
© Boardworks Ltd 2001
Precipitation Reactions – 2nd example
• Most metal hydroxides (except sodium,
potassium and calcium) are insoluble. Reactions
leading to their formation give precipitates.
Copper +
sulphate
aqueous
ammonium
hydroxide
 copper
hydroxide
aqueous
solid
solid
Both reactants
are dissolved (aq).
Copper sulphate
is blue.
+
ammonium
sulphate
aqueous
Copper hydroxide
is insoluble and
precipitates. A pale
blue solid settles at
the bottom of the
test tube.
© Boardworks Ltd 2001
Precipitation Reactions – 3rd example
• Another metal hydroxide that precepitates is
iron(III) hydroxide. Like many transition metals its
compounds are coloured.
Iron + sodium
chloride hydroxide
aqueous
aqueous
Both reactants
are dissolved (aq)
(Iron chloride is
yellow.)
 iron
hydroxide
solid
solid
+
sodium
chloride
aqueous
Iron hydroxide is
insoluble and
precipitates. A
deep brown solid
settles at the
bottom of the test
tube.
© Boardworks Ltd 2001
Precipitation and solubility
•
To work out whether a precipitate will be formed we need
to know the solubility of the compounds that may be
formed. Here are a few general guidelines:
Soluble
Insoluble
All sodium, potassium and
ammonium salts
All nitrates
Most chlorides, bromides and
iodides. (halides)
Silver and lead halides
Most sulphates
Lead, barium and calcium sulphates
Sodium, potassium and
ammonium carbonates
Most carbonates
Sodium, potassium,
ammonium and calcium
hydroxide
Most hydroxides
© Boardworks Ltd 2001
Precipitation and solubility
•
To work out whether a precipitate will be formed when
many ionic compounds react there are four stages:
1.
Write down the names of the
reactants.
Sodium chloride & lead nitrate
2
Write down the ions in the
reactants. (Ignore numbers)
Na+ Cl-
Pb2+ NO3-
Pb2+ Cl-
Na+ NO3-
3
Swap over the + and – ions.
4
Are the products going to be
soluble or insoluble?
Sodium +
chloride
aqueous
lead
nitrate
aqueous
Lead chloride is insoluble so
there will be a precipitate
 lead
chloride
solid
solid
+
sodium
nitrate
aqueous
© Boardworks Ltd 2001
Activity
• Will there be a precipitate if I mix sodium sulphate
and magnesium nitrate?
1.
Write down the names of
the reactants.
2
Write down the ions in the
reactants.
Sodium nitrate &
Na+ SO42-
Magnesium
sulphate
Mg2+ NO3-
3
Swap over the + and – ions.
Mg2+ SO42-
4
Are the products going to
be soluble or insoluble?
Both the products are soluble
there will be no precipitate
Sodium +
sulphate
aqueous
magnesium
nitrate
aqueous
 magnesium
sulphate
aqueous
Na+ NO3-
+ sodium
nitrate
aqueous
© Boardworks Ltd 2001
Activity
• Will there be a precipitate if I mix sodium sulphate
and barium nitrate?
1.
Write down the names of
the reactants.
2
Write down the ions in the
reactants.
Sodium sulphate & barium nitrate
Na+ SO42-
3
Swap over the + and – ions.
4
Are the products going to
be soluble or insoluble?
Sodium +
sulphate
aqueous
barium
nitrate
aqueous
Ba2+ NO3-
Ba2+ SO42-
Na+ NO3-
Barium sulphate is insoluble
so there will be a precipitate
 barium
sulphate
solid
solid
+
sodium
nitrate
aqueous
© Boardworks Ltd 2001
Activity
Separating Precipitates – reminder!
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Neutralisation Reactions
• Acids are substances that:
• Turn litmus red.
• Turn universal indicator yellow, orange
or red.
• Have a pH below 7.
• Form solutions containing H+ ions.
• Bases are substances that:
• Turn litmus blue.
• Turn universal indicator dark green, blue or purple.
• React with the H+ ions in acids.
• Are called alkalis if they dissolve in water.
1
2
3
4
5
Increasingly acid
6
7
8
9
10 11 12 13 14
Increasingly alkali
© Boardworks Ltd 2001
Neutralisation Reactions: Acids
• Common Acids are
Name of acid
Formula
Sulphuric
H2SO4
strong
Hydrochloric
HCl
strong
Nitric
HNO3
strong
CH3COOH
weak
Ethanoic (vinegar)
Strong or Weak?
• Salts
Sulphuric acid
Nitric acid
Hydrochloric acid
Sulphates
Nitrates
Chlorides
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Neutralisation Reactions: Bases
• Common Alkalis are
Name of alkali
Formula
Strong or Weak?
Sodium Hydroxide
NaOH
strong
Potassium Hydroxide
KOH
strong
Calcium Hydroxide
Ca(OH)2
strong
Ammonium
Hydroxide
NH4OH
weak
• Common Bases (neutralise acids but don’t dissolve) are
Type of compound
Contain React with acids to give
Metal Hydroxides
OH-
water + a salt
Metal Oxides
O2-
water + a salt
Metal Carbonates
CO32-
water + a salt + CO2
© Boardworks Ltd 2001
Neutralisation Reactions: acid + base
• A neutralisation reaction is where an acid reacts
with a base to produce a neutral solution of a
salt and water.
Sodium hydroxide
pH 14
neutralisation
Hydrochloric acid
pH 1
1
2
3
Sodium chloride
pH 7
4
5
Increasingly acid
6
7
8
9
10 11 12 13 14
Increasingly alkali
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Neutralisation Naming Salts
• To name the salt formed in a neutralisation:
1 The first part of the name of the salt comes from the
first name of the base
So
Ammonium hydroxide gives ammonium …………
Magnesium oxide gives magnesium …………...
2 The acid gives the last part of the name of the salt.
So
Sulphuric acid make sulphates
Nitric acid makes nitrates
Hydrochloric acid makes chlorides
Eg. Sodium hydroxide + nitric acid forms:
Calcium carbonate + sulphuric acid forms:
Sodium nitrate
calcium sulphate
© Boardworks Ltd 2001
Activity
• Name the salt formed in these neutralisations:
+

Base
Acid
Calcium hydroxide
Hydrochloric acid
Magnesium oxide
Nitric acid
Calcium carbonate
Sulphuric acid
Aluminium
hydroxide
Nitric acid
Potassium hydroxide
Sulphuric acid
Salt?
Calcium chloride
Magnesium nitrate
Calcium sulphate
Aluminium nitrate
Potassium sulphate
© Boardworks Ltd 2001
Neutralisation Reactions: hydroxides
• Each OH- ion reacts with one H+ ion
Reaction with hydroxides:
H+ + OH-  H2O
Eg. Potassium +hydrochloric  water +
hydroxide
acid
KOH
+
HCl
 H2O
Eg. Calcium + sulphuric  water +
hydroxide
acid
Ca(OH)2 +
H2SO4
+
potassium
chloride
KCl
calcium
sulphate
 2H2O +
CaSO4
© Boardworks Ltd 2001
Neutralisation Reactions: oxides
• Neutralisation reactions usually lead to water
being formed.
Reaction with oxides:
2H+ + O2-  H2O
Eg. Calcium + hydrochloric  water +
oxide
acid
CaO
+ 2HCl
 H2O
Eg. Sodium + sulphuric  water +
oxide
acid
Na2O
+
H2SO4
 H2O
calcium
chloride
+
CaCl2
sodium
sulphate
+
Na2SO4
© Boardworks Ltd 2001
Neutralisation Reactions: carbonates
• Each carbonate ion provides one oxygen to join
with two H+ ions. At the same time carbon
dioxide is released
2H+ + CO32-  H2O + CO2
Carbonates:
Eg. Potassium + hydrochloric  water + carbon + potassium
carbonate
acid
dioxide
chloride
K2CO3
Eg. calcium +
carbonate
CaCO3
+ 2HCl
nitric
acid
+ 2HNO3
 H2O + CO2
+
2KCl
 water + carbon + calcium
dioxide nitrate
 H2O
+ CO2
+Ca(NO3)2
© Boardworks Ltd 2001
Neutralisation Equations
• Complete the word equation
Eg. Potassium + hydrochloric 
hydroxide
acid
water
+
Potassium chloride
• Replace the words with the correct formula
KOH
Eg.
+ HCl
 H2O
+
KCl
• Check that it balances (same number of each type of
atom each side).
Eg.
KOH
1*K
+ HCl
Reactants
1*O
2*H

 H2O + KCl
1*Cl
2*H
Products
1*O 1*K
1*Cl
© Boardworks Ltd 2001
Neutralisation Equations
• Complete the word equation
Eg. Magnesium + nitric 
oxide
acid
+
water
Magnesium nitrate
• Replace the words with the correct formula
Eg.
MgO
+ HNO3
 H2O
+
Mg(NO3)2
• Check that it balances (Same number of each type of
atom each side.
Eg.
MgO
1*Mg
+
2 HNO3
Reactants
1*O
1*H
 2 H 2O
1*NO3
2*H
+
Mg(NO3)2
Products
1*O 1*Mg

2*NO3
© Boardworks Ltd 2001
Activity
•
Write balanced equations going through the
same stages as the previous examples
1. word equation,
2. formulae,
3. Balance.
a) sodium hydroxide + hydrochloric acid 
b) magnesium oxide + hydrochloric acid 
c) sodium hydroxide + sulphuric acid 
d) ammonium hydroxide + hydrochloric acid 
e) calcium hydroxide + nitric acid 
© Boardworks Ltd 2001
Activity
• Insoluble salts can be separated by filtering.
• Soluble salts are obtained by evaporating.
Put these in the
correct order
D
A. Check the pH
frequently by testing
drops of the solution.
B. Add the acid slowly
to the alkali.
C. When neutral pour
into the evaporating
basin.
D. Put on safety specs.
B
A
C
F
E
vapour
gauze
Evaporating
basin
tripod
Heat-proof
mat
Bunsen
burner
E. Allow to cool
F. Heat.
© Boardworks Ltd 2001
Redox Reactions:
• Redox is a short way of saying:
• Early on in chemistry these words had
very straightforward meanings
Oxidation meant adding
oxygen to a substance.
Rusting (iron
becoming iron
oxide) is an
example of
oxidation
Reduction
and
oxidation.
Reduction meant taking
oxygen away.
Extracting
iron from iron
oxide in the
blast furnace
is reduction
© Boardworks Ltd 2001
Redox Reactions: Oxidation and ions
• Many redox reactions involve metals and their oxides.
• Whenever metals react with oxygen they form ionic
compounds and the metal loses electrons to form
positively charged ions.
• Eg. When magnesium burns to form magnesium oxide
magnesium atoms (no charge) become magnesium ions
(2+ charge) by losing 2 electrons to oxygen atoms.
Mg
2 e-
O
to give
Mg2+
O2-
Oxidation involves loss of electrons.
© Boardworks Ltd 2001
Redox Reactions:Electron loss.
• Think about what has happened to the magnesium when
it reacts with oxygen.
– It has been oxidised.
– It has lost electrons by changing from Mg  Mg2+
• Magnesium can also lose electrons to things other than
oxygen (e.g. to chlorine or sulphur) and since these also
involve Mg  Mg2+ these too must be oxidation.
O2-
Mg2+
O
Mg
S
Mg2+
S2-
Cl
Oxidation is
loss of
electrons.
Cl- Mg2+ Cl© Boardworks Ltd 2001
Redox Reactions: Electron gain
• Exactly the same reasoning applies to reduction
• Reduction can be the removal of oxygen (e.g. from iron
oxide to form iron or from aluminium oxide in the
electrolysis to extract aluminium.)
• When this happens the metal gets back its electrons.
– Aluminium has been reduced.
– Aluminium has gained electrons
O2Al
Al3+
O2-
Al
Al3+
O2-
Oxygen
removed
O
1
½
O
Reduction
is gain of
electrons.
© Boardworks Ltd 2001
Redox Reactions:Oil Rig
• An easy way of remembering this is “Oil Rig”
O oxidation
I
is
L
loss
of electrons
R Reduction
I
is
G
gain
© Boardworks Ltd 2001
Redox Reactions:Two for one!
• Whenever something is oxidised, something else is
reduced.
• This should be obvious if we use the oil rig definition.
• If something loses electrons – then something else must
have gained them.
•
E.g. When burning magnesium:
– Magnesium loses electrons
(Mg  Mg2+ …..oxidation)
– Oxygen gains electrons
(O  O2-
…….reduction)
•The overall reaction is both Reduction and
Oxidation = Redox
© Boardworks Ltd 2001
Activity
• Say whether the substance in red type is
oxidised or reduced.
Calcium + oxygen

calcium oxide

zinc + water

copper + chlorine

iron + aluminium oxide
oxidised
Zinc oxide + hydrogen
reduced
Copper chloride
reduced
Aluminium + iron oxide
oxidised
© Boardworks Ltd 2001
Activity • If the first substance is oxidised, what has
been reduced or vice versa. (Use whichever
definition of oxidation and reduction seems
easier to apply.).
Calcium + oxygen

calcium oxide
oxidised
Oxygen is reduced. Each oxygen atom gains 2 e-
Zinc oxide + hydrogen
reduced

copper + chlorine
Chlorine is Oxidised. It gains an electron Cl-  ½Cl2
Aluminium + iron oxide
oxidised
zinc + water
Hydrogen is oxidised. It gains oxygen
Copper chloride
reduced


iron + aluminium oxide
Iron is reduced. It loses oxygen
© Boardworks Ltd 2001
Activity
• Across:
5 tells us whether acid or
alkali
11 reaction of an acid with a
base
•
1
2
3
4
6
7
8
9
10
Down
a solid forms in a solution
loss of electrons
competition reaction
gives solutions containing
H+ ions
to break down into smaller
particles
removal of oxygen
state of balance
soluble base
ionic compound formed in
neutralisations
© Boardworks Ltd 2001
Activity
• Match them up
Thermal decomposition
Endothermic
Dehydrating copper sulphate
A solid forms within a solution
Metal displacement
Reversible reaction
Precipitation
A salt and water is formed
Alkali
Reaction in a state of balance
Neutralisation
Oxidation
Reduction
Thermit reaction
Removal of oxygen
Breaking up with heat
Soluble base
Equilibrium
Takes in energy – gets cold
Loss of electrons
© Boardworks Ltd 2001
When heated the orange powder erupted like a
volcano producing a huge pile of green powder
that had less mass than the orange material. Is
this?
1.
2.
3.
4.
Neutralisation
Thermal decomposition
Displacement
Precipitation
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When the two colourless solutions mixed a
yellow solid formed which sank to the bottom of
the test tube Is this?
1.
2.
3.
4.
Neutralisation
Thermal decomposition
Displacement
Precipitation
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When the copper was placed in the silver nitrate
solution snow like crystals of silver seemed to
grow out from the copper. Is this?
1.
2.
3.
4.
Equilibrium
Thermal decomposition
Displacement
Precipitation
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When the washing soda was added to the
lemon juice it fizzed and the pH rose towards 7.
Is this?
1.
2.
3.
4.
Neutralisation
Thermal decomposition
Displacement
Oxidation
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Which of the oxides shown will thermally
decompose most easily?
1.
2.
3.
4.
Mercury oxide
Potassium oxide
Iron oxide
Gold oxide
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Which of the salts below might be formed when
nitric acid neutralises a metal hydroxide?
1.
2.
3.
4.
Potassium hydroxide
Potassium nitrate
Ammonium nitrate
Calcium sulphate
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Which of the mixtures below will result in a
metal displacement reaction?
1.Potassium oxide and gold
2.Magnesium and sodium nitrate
3.Copper and gold nitrate
4.Aluminium and calcium sulphate
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Which of the mixtures below will result in a nonmetal displacement reaction?
1.Potassium chloride and iodine
2.Potassium bromide and iodine
3.Potassium fluoride and chlorine
4.Potassium iodide and chlorine
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Which of the elements in red (below) is oxidised
in the reaction? (Oil Rig!)?
1.Ca +
2.2Li +
3.2Al +
4.HNO3
CuO  CaO + Cu
2HCl  2LiCl + H2
Fe2O3  Al2O3 + 2Fe
+ CuO  CuNO3 + H2O
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Which compound can you be sure is soluble in
water?
1.
2.
3.
4.
Manganese nitrate
Osmium iodide
Thallium chloride
Palladium sulphate
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