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© Boardworks Ltd 20051 of 69
KS4 Chemistry
Chemical Reactions
© Boardworks Ltd 20052 of 69
Chemical Reactions
Displacement and precipitation
Neutralization
Introducing chemical reactions
Redox
Thermal decomposition
Summary activities
Contents
© Boardworks Ltd 20053 of 69
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).
reactants productschemical reaction
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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oxidation &reduction
neutralization
precipitation
reversible
displacement:metals
exothermic &endothermic
thermaldecomposition
displacement:non-metals
Types of chemical reaction
chemical reaction
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Exothermic and endothermic reactions
What are exothermic and endothermic reactions?
exothermic reactions give out energy – they get hot
endothermic reactions take in energy – they get cold
ex = out (as in ‘exit’)
en = in (as in ‘entrance’)
Most chemical reactions are exothermic.
thermic = relating to heat
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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
Examples of exothermic reactions
Many exothermic reactions occur in the lab and in everyday life. Can you think of six exothermic reactions?
Exothermic reactions
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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.
In equations for irreversible reactions, reactants and products are joined by a ‘one-way’ arrow.
hydrochloricacid
magnesiumchloride
magnesium + hydrogen+
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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.
anhydrous copper sulfate
hydratedcopper 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:
hydrogennitrogen + ammonia
N2 (g) 3H2 (g) 2NH3 (g)+
No matter how long the reaction is left, there will always bea mixture of nitrogen, hydrogen and ammonia.
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Chemical Reactions
Displacement and precipitation
Neutralization
Introducing chemical reactions
Redox
Thermal decomposition
Summary activities
Contents Contents
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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.
potassium carbonate: is not thermallydecomposed
calcium carbonate: decomposes onstrong heating
silver carbonate: decomposes ongentle heating
increase in reactivity of metal
For example:
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easy
hard
medium
easy
medium
How easy will these metal compounds be to decompose: easy, medium or difficult?
Thermal decomposition – easy or hard?
mercury oxide
sodium oxide
iron oxide
silver oxide
zinc oxide
Compound Decomposition
incr
ease
in
rea
ctiv
ity
potassiumsodiumcalcium
magnesiumaluminium
zinciron
coppermercury
silvergold
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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):
Quicklime is used to make concrete and calcium hydroxide (slaked lime).
CaOCaCO3 + CO2
calciumoxide
calciumcarbonate + carbon
dioxide
heat
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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 condenses at the top of the test tube, where it is cooler
Thermal decomposition of metal oxides
2Hg2HgO + O2
oxygen gas escapes
mercurymercury
oxide + oxygenheat
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Chemical Reactions
Displacement and precipitation
Neutralization
Introducing chemical reactions
Redox
Thermal decomposition
Summary activities
Contents
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The reactivity series
incr
ease
in
rea
ctiv
ity
potassiumsodiumcalcium
magnesiumaluminium
zinciron
coppermercury
silvergold
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.
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.
+less
reactive metal
compound of more reactive
metal
compound of less
reactive metal
more reactive
metal+
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Displacement reactions – examples
Will silver react with magnesium chloride?
Magnesium is more reactive than copper, so it displaces copper from its compound.
Silver is less reactive than magnesium, so it does not displace magnesium from its compound.
Will magnesium react with copper chloride?
+magnesiumcopper
chloride +magnesium
chloridecopper
+silvermagnesium
chloride no reaction
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The Thermit process
Aluminium is more reactive than iron and displaces it from the oxide.
The Thermit process is a displacement reaction between aluminium and iron (III) oxide.
aluminiumoxide
+ ironiron oxidealuminium +
Fe2O3 Al2O3Al + Fe+
magnesiumfuse
aluminium powder (thermite)
iron oxide
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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.
The Thermit process
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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.
F2 (aq) 2NaCl (aq) 2NaF (aq) Cl2 (aq)++
The less reactive halogen is displaced from the compound and the more reactive halogen bonds with the metal to form a new metal halide.
I
Br
Cl
F
decrease in reactivity
For example:
sodiumchloride
sodiumfluoride
chlorinefluorine + +
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Displacement reactions of halogens
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Is there a displacement reaction?
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What are the symbols for these physical states?
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.
solid
liquid
gas
aqueous(dissolved in water)
(g)
(l)
(g)
(aq)
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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.
hydrochloricacid
sodiumchloride
sulfursodium
thiosulfate + + watersulfur
dioxide ++
Na2S2O3
(aq)2HCl(aq)
2NaCl(aq)
S(s)++
SO2
(g)H2O(l)
+ +
Both reactants are colourless.
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:
ammoniumhydroxide
copper (II)hydroxide
copper (II)sulfate + +
ammoniumsulfate
CuSO4
(aq)2NH4OH
(aq)Cu(OH)2
(s)++
(NH4)2SO4
(aq)
Copper (II) sulfate solution is blue.
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.
sodiumhydroxide
iron (III)hydroxide
iron (III)chloride + +
sodiumchloride
FeCl3 (aq) 3NaOH (aq) Fe(OH)3 (s) ++ 3NaCl (aq)
Iron (III) chloride solution is yellow.
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.
All compounds of sodium, potassium and ammonium.
All nitrates.
Most chlorides, except silver and lead chlorides.
Most sulfates, except lead, barium and calcium sulfates.
Most carbonates and hydroxides, except those of sodium, potassium and ammonium.
Soluble
Insoluble
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There are three steps to working out whether a precipitate will be formed in a reaction:
1. Write down the names of the reactants.
3. Are the products soluble or insoluble?
Predicting precipitation reactions
sodium chloride & lead nitrate
Lead chloride is insoluble and will form a precipitate.
2. Swap over the non-metal.
sodium nitrate & lead chloride
Example
leadnitrate
sodiumnitrate
sodiumchloride + +
leadchloride
2NaCl (aq) Pb(NO3)2 (aq) 2NaNO3 (aq) ++ PbCl2 (s)
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Will a precipitate be formed when sodium nitrate and magnesium sulfate react?
Will there be a precipitation reaction? (1)
1. Write down the names of the reactants.
3. Are the products soluble or insoluble?
sodium nitrate &magnesium sulfate
Both products are soluble so no precipitate will form.
2. Swap over the non-metal.
sodium sulfate & magnesium nitrate
magnesiumsulfate
sodiumsulfate
sodiumnitrate + +
magnesiumnitrate
2NaNO3 (aq) MgSO4 (aq) Na2SO3 (aq) ++ 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.
3. Are the products soluble or insoluble?
sodium sulfate &barium nitrate
Barium sulfate is insoluble and will form a precipitate.
2. Swap over the non-metal.
sodium nitrate & barium sulfate
bariumnitrate
sodiumnitrate
sodiumsulfate + +
bariumsulfate
Na2SO4 (aq) Ba(NO3)2 (aq) 2NaNO3 (aq) ++ BaSO4 (s)
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Precipitation: true or false?
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Chemical Reactions
Displacement and precipitation
Neutralization
Introducing chemical reactions
Redox
Thermal decomposition
Summary activities
Contents
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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.
weakacid
strongacid
1 2 3 4 5 6 7 8 9 10 11 12 13 14
neutral
Universal indicator pH scale
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HNO3
Common acids
What are some common acids?
sulfuric acid
hydrochloric acid
nitric acid
ethanoic acid (vinegar)
H2SO4
HCl
CH3COOH
strong
strong
strong
weak
Acid Formula Strength
1 2 3 4 5 6 7 8 9 10 11 12 13 14
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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.
weak base
strongbase
1 2 3 4 5 6 7 8 9 10 11 12 13 14
neutral
Universal indicator pH scale
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basespH > 7
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.
Most carbonates and hydroxides are insoluble in water, except those of sodium, potassium and ammonium.
alkalissolublebases
All alkaline solutions contain OH– ions. The more OH– ions in the solution, the stronger the alkali.
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Common bases
Ca(OH)2
What are some common bases?
sodium hydroxide *
potassium hydroxide *
calcium hydroxide *
ammonia *
NaOH
KOH
NH3
strong
strong
strong
weak
1 2 3 4 5 6 7 8 9 10 11 12 13 14
calcium carbonate CaCO3 weak
Base Formula Strength
* = the base is also an alkali
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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.
alkali saltacid + water+
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The first part of the salt comes from the first part of the base, for example:
The second part of the salt comes from the acid:
Naming salts
The salt formed in a neutralization reaction takes its name from both the base and the acid.
magnesium oxide magnesium salt ammonium hydroxide ammonium salt
sulfuric acidsulfate
hydrochloric acid chloride
nitric acid nitrate
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sulfuric acid potassium sulfatepotassium hydroxide
calcium chloride
magnesium nitrate
copper sulfate
aluminium nitrate
What is the name of the salt?
Base Acid Salt
calcium hydroxide
magnesium oxide
copper oxide
aluminium hydroxide
hydrochloric acid
nitric acid
nitric acid
sulfuric acid
What are the names of the salts formed from these bases and acids?
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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:
OH– H2OH+ +
hydrochloricacid
potassiumchloride
potassiumhydroxide + + water
KOH (aq) HCl (aq) KCl (aq) ++ H2O (aq)
For example:
Ca(OH)2 (aq) H2SO4 (aq) CaSO4(aq) ++ 2H2O (aq)
sulfuricacid
calciumsulfate
calciumhydroxide + + water
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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:
O2– H2O2H+ +
hydrochloricacid
calciumchloride
calciumoxide + + water
CaO (aq) 2HCl (aq) CaCl2 (aq) ++ H2O (aq)
For example:
sulfuricacid
coppersulfate
copperoxide + + water
CuO (aq) H2SO4 (aq) CuSO4 (aq) ++ 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:
CO32– H2O2H+ + + CO2
nitricacid
calciumnitrate
calciumcarbonate + + water +
carbondioxide
CaCO3 (aq)
2HNO3
(aq)
Ca(NO3)2
(aq)++ H2O
(aq)+ CO2
(g)
For example:
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Complete the neutralization reaction
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Obtaining insoluble salts
How can insoluble salts be obtained?
2. The mixture of products is filtered, trapping the salt.
1. The acid and alkali are mixed, and the salt forms by precipitation.
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.
sulfuricacid
bariumsulfate
bariumhydroxide + + water
BaOH2 (aq) H2SO4 (aq) BaSO4 (s) ++ 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)?
2. More copper oxide is added until no more will dissolve. This means that all the acid has been used up.
1. Copper oxide is added to sulfuric acid. When it is heated, the copper oxide dissolves, forming a blue solution.
3. The mixture is filtered to remove the excess copper oxide.
For example, obtaining copper sulfate from copper oxide and sulfuric acid:
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)?
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.
hydrochloricacid
sodiumchloride
sodiumhydroxide + + water
NaOH (aq) HCl (aq) NaCl (aq) ++ H2O (aq)
For example, the reaction between sodium hydroxide and hydrochloric acid produces sodium chloride, which is soluble.
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Obtaining soluble salts from alkalis (2)
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.
To run the titration:
5. The salt solution is heated to evaporate the water. Crystals of sodium chloride will remain.
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.
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Matching reactants and salts
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Neutralization: true or false?
© Boardworks Ltd 200554 of 69
Chemical Reactions
Displacement and precipitation
Neutralization
Introducing chemical reactions
Redox
Thermal decomposition
Summary activities
Contents
© Boardworks Ltd 200555 of 69
Oxidation = adding oxygen to a substance.
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.
Reduction = removing oxygen from a substance.
E.g. production of iron oxide during rusting
E.g. extracting iron from ironoxide 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?
For example, when magnesium burns to form magnesium oxide, each magnesium atom loses 2 electrons and becomes a magnesium ion with a +2 charge.
Metal atoms lose electrons to form positive ions. The oxygen atoms accept these electrons and form negative ions.
O2-Mg Mg2+
2 electrons
O
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Oxidation and electron loss
What has happened to magnesium when it reacts with oxygen?
It has been oxidized.
It has lost electrons by changing from Mg to Mg2+.
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.
Oxidation is the loss of electrons.
O2-Mg Mg2+
2 electrons
O
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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.
O
O
O
Fe
Fe
2 electrons from each ion
O2-
O2-
O2-
Fe3+
Fe3+
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Reduction and electron gain
What has happened to iron when oxygen is removed?
It has been reduced.
It has gained electrons by changing from Fe3+ to Fe.
Reduction is the gain of electrons.
O
O
O
Fe
Fe
2 electrons from each ion
O2-
O2-
O2-
Fe3+
Fe3+
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Electrons and OILRIG
An easy way to remember what happens to electrons during oxidation and reduction is to think…
O
I
L
R
I
G
xidation
s
oss
eduction
s
ain
…OILRIG
…of electrons
…of electrons
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The overall reaction is reduction and oxidation = redox.
magnesium loses electrons:
Redox reactions
When a substance is oxidized, it loses electrons. Another substance must gain these electrons and become reduced.
For example, when magnesium burns:
Oxidation and reduction always take place together.
oxygen gains electrons:
Mg Mg2+ = oxidation
O O2– = reduction
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oxidized
reduced
Oxidized or reduced?
For each reaction, decide which product is oxidized and which product is reduced.
reduced
reducedoxidized
oxidized
calcium + oxygen calcium oxide
zinc oxide hydrogen+ + zinc water
aluminium iron oxide +iron aluminium oxide+
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Oxidized or reduced?
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Chemical Reactions
Displacement and precipitation
Neutralization
Introducing chemical reactions
Redox
Thermal decomposition
Summary activities
Contents
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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