Ks4 chemical reactions

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oxidation
and reduction
neutralisation
precipitation reversible
reactions
displacement
reactions:
metals
exothermic
and endothermic
thermal
decomposition
displacement
reactions:
non-metals
Types of
chemical change

Thermal decomposition
• 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.
For example:
Potassium carbonate is not thermally decomposed.
Calcium carbonate decomposes on strong heating
Silver carbonate decomposes on gentle heating
Getsharder

Generally, the more reactive the metal, the
more difficult it is to decompose its
compounds.
Fill in the last column: easy, medium or hard.
Potassium
sodium
calcium
magnesium
aluminium
zinc
iron
copper
mercury
silver
gold
Increasing
Compound How easy to decompose
Mercury oxide
Sodium oxide
Iron oxide
Silver oxide
Zinc oxide
easy
hard
medium
easy
medium
Thermal decomposition

Thermal decomposition of carbonates
• 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).
1500°C
limestone
Hot air
calcium oxide (lime)
waste
air and
carbon
dioxide
Calcium
Carbonate
Calcium
oxide
Carbon
dioxide
+

Thermal decomposition of metal oxides
• Most metal oxides are
thermally stable (i.e. do not
decompose when heated).
• Oxides of the least reactive
metals can be thermally
decomposed more easily.
• For example, silver oxide
begins to break up at about
160o
C and mercury oxide
decomposes when heated
strongly.
Mercury Oxide Mercury oxygen+
Hg
Hg
Hg
Hg
O O
O
O
Heat
HgHg
Hg HgOO
O O
HgHg
Hg HgOO
O O
HgHg
Hg HgOO
O O
HgHg
Hg HgOO
O O
mercury oxide
decomposes
mercury metal
and oxygen
formed

Exothermic and endothermic reactions
• Exothermic reactions give out heat (gets hot).
• Endothermic reactions take in heat (gets 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
Ex = out (as in exit)Ex = out (as in exit)
En = in (as in entrance)En = in (as in entrance)

• 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
Exothermic and endothermic reactions

Displacement reactions: metals
• 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
displaceddisplaced and so is present as
the metal at the end of the
reaction.
Potassium
sodium
calcium
magnesium
aluminium
zinc
iron
copper
silver
gold
Increasingreactivity
A more reactive metal (higher in the reactivity series) will
displace a less reactive metal from its compound.

• Copper is quite low in the activity series.
• Several metals will displace it from its compounds.
magnesium copper
sulphate
solution
magnesium
sulphate solution
copper metal
Magnesium + Copper
sulphate
 Magnesium
sulphate
+ Copper
more
reactive
less
reactive
Magnesium wins the competition.
Copper is displaced.
K
Na
Ca
Mg
Al
Zn
Fe
Cu
Ag
Au

Here are some actual photos.
The colour changes from blue to red/black as
copper metal is displaced.
Magnesium + Copper
sulphate
 Magnesium
sulphate
+ Copper
more
reactive
less
reactive
Magnesium wins the
competition. Copper is displaced
photograph
at end of
reaction
photograph
at start of
reactionK
Na
Ca
Mg
Al
Zn
Fe
Cu
Ag
Au

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 broken railway tracks.
Aluminium + Iron
Oxide
 Aluminium
Oxide
+ Iron
more
reactive
less
reactive
Aluminium wins the competition.
Iron is displaced and melts at the
high temperatures produced.
K
Na
Ca
Mg
Al
Zn
Fe
Cu
Ag
Au
iron oxide +
aluminium
powder
magnesium
fuse

Here is a photo of the thermit reaction being
carried out in a laboratory.
iron oxide +
aluminium
powder
magnesium
fuse

Predict which mixtures will result in a reaction.
Metal 
Solution
Iron Magnesium Zinc Copper
Iron
chloride
Magnesium
nitrate
Zinc nitrate
Copper
sulphate
Yes Yes No
No No No
No Yes No
Yes Yes Yes

L.O. Be able to describe some
properties of elements in groups1, 7
and 8

TASKS
•List all the elements that have 7 and 8
electrons in the outer shell
•Write two halogen displacement reactions
and state colour of final solution
•Draw the atomic configuration of
Helium, Neon and Chlorine and Lithium
•State some of the properties of the
halogens, noble gases and alkali metals
105

Displacement reactions: halogens
• 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.
Increasingreactivity
Fluorine
Chlorine
Bromine
Iodine
• We can often tell which halogen is present from the
colour of the solution.

For example, if chlorine solution is added to
sodium bromide.
sodium
bromide
solution
sodium chloride
solution
bromine
Chlorine + Sodium
Bromide
 Sodium
Chloride
+ Bromine
more
reactive
less
reactive
Chlorine wins the competition.
Bromine (red) is displaced.
F
Cl
Br
I
At
chlorine
solution

The compounds of the halogens with Group 1 metals
are all colourless.
Halogen
Halide 
Chlorine
solution
Bromine
solution
Iodine
Solution
Potassium
chloride
Potassium
bromide
Potassium
Iodide
Br2 I2
Br2
I2
I2 I2
Predict what colour these will be after mixing.

Chlorine +
When writing equations for halogen displacement
reactions you must remember that – when in the
form of the element – halogens exist in pairs.
For chlorine and sodium bromide:
+ bromine
sodium
chloride
Sodium
bromide
Cl2(aq) + 2NaBr(aq) 2NaCl(aq) + Br2(aq)
F
Cl
Br
I
At
Cl More
reactive
Br Less
reactive
Solution goes
yellow/brown as
bromine is produced.

• 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.
F
Cl
Br
I
At
1) iodine + sodium bromide solution
2) bromine + sodium chloride solution
3) chlorine + sodium iodide solution
No reaction
No reaction
sodium chloride + iodine
Cl2(g) + 2NaI(aq) 2NaCl(aq) + I2(aq)
Predict whether or not a chemical reaction will
occur.

Reversible and irreversible reaction
• Most chemical reactions are considered
irreversible in that the new products are not
readily changed back into reactants. For
example, 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 “one-way arrow.”
magnesium + hydrochloric magnesium + hydrogen
acid chloride

• 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
hydrated copper
sulphate
Heat
anhydrous copper
sulphate
steam
CuSO4.5H20 CuSO4 + 5H2O
these decompose these combine
Reversible reactions

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.

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)  2NH3 (g) (exothermic)
Low temperature High pressure

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) solid
–(l) liquid
–(g) gas
–(aq) aqueous (dissolved in water)

A precipitation reaction that is often used to measure
reaction rates occurs between sodium thiosulphate and
hydrochloric acid.
Sodium + hydrochloric sodium + sulphur + water + sulphur
thiosulphate acid chloride dioxide
Both reactants are
colourless and
dissolved (aq)
Sulphur is insoluble
and precipitates.
This makes the
solution go cloudy.
aqueous aqueous aqueous solid liquid gassolid
Precipitation reactions – first example

Most metal hydroxides (except sodium, potassium and
calcium) are insoluble. Reactions leading to their formation
give precipitates.
Copper + ammonium copper + ammonium
sulphate hydroxide hydroxide sulphate
aqueous aqueous solid aqueoussolid
Copper hydroxide is
insoluble and
precipitates. A pale
blue solid settles at
the bottom of the
test tube.
Both reactants
are dissolved (aq).
Copper sulphate is
blue.
Precipitation reactions – second example

Another metal hydroxide that precipitates is iron(III)
hydroxide. Like many transition metals its compounds are
coloured.
Iron + sodium iron + sodium
chloride hydroxide hydroxide chloride
Iron hydroxide is
insoluble and
precipitates. A deep
brown solid settles
at the bottom of the
test tube.
Both reactants
are dissolved (aq)
(iron chloride is
yellow).
Precipitation reactions – third example

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 nitratesnitrates
Most chlorides, bromides and
iodides. (halides)
Silver and lead halides
Most sulphatessulphates Lead, barium and calcium sulphates
Sodium, potassium and
ammonium carbonatescarbonates
Most carbonates
Sodium, potassium,
ammonium and calcium
hydroxidehydroxide
Most hydroxides

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)
3 Swap over the + and – ions.
4 Are the products going to be
soluble or insoluble?
Na+
Cl-
Pb2+
NO3
-
Pb2+
Cl-
Na+
NO3
-
Lead chloride is insoluble so
there will be a precipitate
Sodium + lead lead + sodium
chloride nitrate chloride nitrate
Precipitation and solubility

Will there be a precipitate if I mix sodium
sulphate and magnesium nitrate?
Sodium nitrate & Magnesium
sulphate
1 Write down the names of the
reactants.
reactants.
Na+
SO4
2-
Mg2+
NO3
-
Mg2+
SO4
2-
Na+
NO3
-
Both the products are soluble
there will be no precipitate.
Sodium + magnesium magnesium + sodium
sulphate nitrate sulphate nitrate
aqueous aqueous aqueous aqueous

Will there be a precipitate if I mix
sodium sulphate and barium nitrate?
Sodium sulphate & barium nitrate1 Write down the names of the
reactants.
reactants.
Na+
SO4
2-
Ba2+
NO3
-
Ba2+
SO4
2-
Na+
NO3
-
Barium sulphate is insoluble so
there will be a precipitate.
Sodium + barium barium + sodium
sulphate nitrate sulphate nitrate

Separating Precipitates – reminder!

Neutralisation reactions
• AcidsAcids are substances that:
• Turn litmus red.
• Turn universal indicator yellow, orange
or red.
• Have a pH below 7.
• Form solutions containing H+
ions.
• BasesBases 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 14131211109876543
Increasingly acid Increasingly alkali

Neutralisation reactions: acids
• Common AcidsCommon Acids are
Name of acid Formula Strong or Weak?
Sulphuric
Hydrochloric
Nitric
Ethanoic (vinegar)
H2SO4
HCl
HNO3
CH3COOH
strong
strong
strong
weak
• SaltsSalts
Sulphuric acid
Sulphates
Nitric acid
Nitrates Chlorides
Hydrochloric acid

Neutralisation reactions: bases
• Common alkalisCommon alkalis are
Name of alkali Formula Strong or Weak?
Sodium Hydroxide
Potassium Hydroxide
Calcium Hydroxide
Ammonium Hydroxide
NaOH
KOH
Ca(OH)2
NH4OH
strong
strong
strong
weak
• Common basesCommon bases (neutralise acids but don’t dissolve) are
Type of compound Contain React with acids to give
Metal Hydroxides
Metal Oxides
Metal Carbonates
OH-
O2-
CO3
2-
water + a salt
water + a salt
water + a salt + CO2

Neutralisation reactions: acid + base
A neutralisation reaction is where an acidacid reacts with a
basebase to produce a neutral solution of aa saltsalt and waterwater.
1 2 14131211109876543
Increasingly acid Increasingly alkali
sodium hydroxide
pH 14
hydrochloric acid
pH 1
neutralisation
sodium chloride
pH 7

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 sulphatessulphates
Nitric acid makes nitratesnitrates
Hydrochloric acid makes chlorideschlorides
Eg. Sodium hydroxide + nitric acid forms:
Calcium carbonate + sulphuric acid forms:
Sodium nitrate
calcium sulphate

Name the salt formed in these neutralisations:
Base Acid Salt?
Calcium hydroxide Hydrochloric acid
Magnesium oxide Nitric acid
Calcium carbonate Sulphuric acid
Aluminium
hydroxide
Nitric acid
Potassium hydroxide Sulphuric acid
Calcium chloride
Magnesium nitrate
Calcium sulphate
Aluminium nitrate
Potassium sulphate
+ 

Neutralisation reactions: hydroxides
Each OH-
ion reacts with one H+
ion.
Reaction with hydroxides: H+
+ OH-
 H2O
Eg. Potassium +hydrochloric  water + potassium
hydroxide acid chloride
KOHOH + HHCl  HH22OO + KCl
Eg. Calcium + sulphuric  water + calcium
hydroxide acid sulphate
Ca(OHOH)22 + HH22SO4  2H2H22OO + CaSO4

Neutralisation Reactions: oxides
Neutralisation reactions usually lead to water being formed.
Reaction with oxides: 2H+
+ O2-
 H2O
Eg. Calcium + hydrochloric  water + calcium
oxide acid chloride
CaOO + 2H2HCl  HH22OO + CaCl2
Eg. Sodium + sulphuric  water + sodium
oxide acid sulphate
Na2OO + HH22SO4  HH22OO + Na2SO4

Neutralisation Reactions:
carbonates
Each carbonate ion provides one oxygen to join with two
H+ ions. At the same time carbon dioxide is released.
Carbonates: 2H+
+ CO3
2-
 H2O + CO2
Eg. Potassium + hydrochloric  water + carbon + potassium
carbonate acid dioxide chloride
K2COCO33 + 2H2HCl  HH22OO + COCO22 + 2KCl
Eg. calcium + nitric  water + carbon + calcium
carbonate acid dioxide nitrate
CaCOCO33 + 2H2HNO3  HH22OO + CO2 +Ca(NO3)2

Neutralisation equations
Eg. Potassium + hydrochloric  +
hydroxide acid
Complete the word equation
Eg. KOH + HHCl  +
water Potassium chloride
HH22OO KCl
Replace the words with the correct formula
Check that it balancesbalances (same number of each type of
atom each side).
Eg. KOH + HHCl  HH22OO + KCl
Reactants
1*K 1*O 2*H 1*Cl
Products
2*H 1*O 1*K 1*Cl


Eg. Magnesium + nitric  +
oxide acid
Complete the word equation
Eg. MgO + HHNO3  +
water Magnesium nitrate
HH22OO Mg(NO3)2
Replace the words with the correct formula
Check that it balancesbalances (Same number of each type of
atom each side.
Reactants
1*Mg 1*O 1*H1*H 1*NO1*NO33
Products
2*H2*H 1*O 1*Mg 2*NO2*NO33
Eg. MgO + HHNO3  H2O + Mg(NO3)2
2 2

Neutralisation equations

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

• Insoluble salts can be separated by filtering.
• Soluble salts are obtained by evaporating.
bunsen
burner
evaporating
basingauze
tripod
heat-proof
mat
vapour
Put these in the
correct order.
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.
E. Allow to cool
F. Heat.
D B A C F E

Redox Reactions
• Redox is a short way of saying:
Reduction
and
oxidation
Oxidation meant addingOxidation meant adding
oxygen to a substance.oxygen to a substance.
Rusting (iron
becoming iron
oxide) is an
example of
oxidation.
Reduction meant takingReduction meant taking
oxygen away.oxygen away.
Extracting
iron from iron
oxide in the
blast furnace
is reduction.
• Early on in chemistry these words had
very straightforward meanings.

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 O2 e-
to give Mg2+
O2-
Oxidation involves loss of electrons.

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.
Mg Oxidation is
the loss of
electrons.
Mg2+
S2-
S
Mg2+
O2-
O
Mg2+
Cl-Cl-
Cl
Mg
Mg2+
S2-
S
Mg2+
O2-
O
Mg2+
Cl-Cl-
Cl

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
Al3+
O2-
O2-
O2-
Al3+
Oxygen
removed
Reduction
is the gain of
electrons.
Al
Al
O
O
1
½

Redox Reactions: oil rig
An easy way of remembering this is “Oil RigOil Rig”!
O oxidationO oxidation
II isis
L lossL loss
R reductionR reduction
II isis
G gainG gain
of electrons

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.
• For example, when burning magnesium:
– Magnesium loses electrons
(Mg Mg2+
…..oxidation)
– Oxygen gains electrons
(O O2-
…….reduction)
The overall reaction is both
RedReduction and OxOxidation = RedoxRedox

Say whether the substance in red type is
oxidised or reduced.
CalciumCalcium + oxygen calcium oxide
ZincZinc oxide + hydrogen zinc + water
CopperCopper chloride copper + chlorine
AluminiumAluminium + iron oxide iron + aluminium oxide
oxidised
reduced
reduced
oxidised

If the first substance is oxidised, what has been
reduced or vice versa (use whichever definition of
oxidation and reduction seems easier to apply).
CalciumCalcium + oxygen calcium oxide
ZincZinc oxide + hydrogen zinc + water
CopperCopper chloride copper + chlorine
AluminiumAluminium + iron oxide iron + aluminium oxide
oxidised
reduced
reduced
oxidised
Oxygen is reduced. Each oxygen atom gains 2 e-.
Hydrogen is oxidised. It gains oxygen.
Chlorine is oxidised. It gains an electron Cl-
½Cl2
Iron is reduced. It loses oxygen.

• Across:
5 tells us whether acid or
alkali
11 reaction of an acid with a
base
• Down
1 a solid forms in a solution
2 loss of electrons
3 competition reaction
4 gives solutions containing
H+ ions
6 to break down into smaller
particles
7 removal of oxygen
8 state of balance
9 soluble base
10 ionic compound formed in
neutralisations

Match them up
Thermal decomposition Dehydrating copper sulphate
Endothermic A solid forms within a solution
Metal displacement A salt and water is formed
Reversible reaction Alkali
Precipitation Reaction in a state of balance
Neutralisation Thermit reaction
Oxidation Removal of oxygen
Reduction Breaking up with heat
Soluble base Takes in energy – gets cold
Equilibrium Loss of electrons

When heated the orange powder erupted like a
volcano producing a huge pile of green powder
that had less mass than the orange material.
What type of reaction is this?
1. Neutralisation
2. Thermal decomposition
3. Displacement
4. Precipitation

When the two colourless solutions mixed a
yellow solid formed which sank to the bottom of
the test tube. What type of reaction is this?
1. Neutralisation
3. Displacement
4. Precipitation

When the copper was placed in the silver nitrate
solution snow-like crystals of silver seemed to
grow out from the copper.
1. Equilibrium
3. Displacement
4. Precipitation

When the washing soda was added to the
lemon juice it fizzed and the pH rose towards 7.
1. Neutralisation
3. Displacement
4. Oxidation

Which of the oxides shown will thermally
decompose most easily?
1. Mercury oxide
2. Potassium oxide
3. Iron oxide
4. Silver oxide

Which of the salts below might be formed when
nitric acid neutralises a metal hydroxide?
1. Potassium hydroxide
2. Potassium nitrate
3. Ammonium nitrate
4. Calcium sulphate

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 silver nitrate
4.Aluminium and calcium sulphate

Which of the mixtures below will result in a
non-metal displacement reaction?
1.Potassium chloride and iodine
2.Potassium bromide and iodine
3.Potassium fluoride and chlorine
4.Potassium iodide and chlorine

Which of the elements in red (below) is oxidised
in the reaction? (Oil Rig!)
1.Ca + CuCuO  CaO + Cu
2.2Li + 2HHCl  2LiCl + H2
3.2AlAl + Fe2O3  Al2O3 + 2Fe
4.HNO3 + CuCuO  CuNO3 + H2O

Which compound can you be sure is soluble
in water?
1. Manganese nitrate
2. Osmium iodide
3. Thallium chloride
4. Palladium sulphate

Ks4 chemical reactions

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