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© Boardworks Ltd 2003

Chemical Reactions


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oxidationand reduction

neutralisation

precipitation reversiblereactions

displacementreactions:

metals

exothermicand endothermic

thermaldecomposition

displacementreactions:

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 Get

s h

ard

er


Generally, the more reactive the metal, the more difficult it is to decompose its compounds.

Fill in the last column: easy, medium or hard.

Potassiumsodiumcalcium

magnesiumaluminium

zinciron

coppermercury

silvergold

Incr

ea

sing

re

activ

ity

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)

wasteair and carbondioxide

Calcium Carbonate

Calcium oxide

Carbondioxide

+


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 160oC and mercury oxide decomposes when heated strongly.

Mercury Oxide Mercury oxygen+

Hg

HgHg

Hg

O OO

O

Heat

HgHg

Hg HgOO

O OHgHg

Hg HgOO

O O

HgHg

Hg HgOO

O OHgHg

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.

Potassiumsodiumcalcium

magnesiumaluminium

zinciron

coppersilvergold

Incr

ea

sing

re

activ

ity

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 coppersulphatesolution

magnesiumsulphate solution

copper metal

Magnesium + Copper sulphate

Magnesium sulphate

+ Copper

more reactive

less reactive

Magnesium wins the competition. Copper is displaced.

KNaCaMgAlZnFeCuAgAu



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

NaCaMgAlZnFeCuAgAu



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.

KNaCaMgAlZnFeCuAgAu


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


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.

Incr

ea

sing

re

activ

ity

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 chloridesolution

bromine

Chlorine + Sodium Bromide

Sodium Chloride

+ Bromine

more reactive

less reactive

Chlorine wins the competition. Bromine (red) is displaced.

FClBrI

At

chlorine solution



The compounds of the halogens with Group 1 metals are all colourless.

HalogenHalide

Chlorine

solution

Bromine solution

Iodine Solution

Potassium chloride

Potassium bromide

Potassium Iodide

Br2 I2

Br2I2

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:

+ brominesodium chloride

Sodiumbromide

Cl2(aq) + 2NaBr(aq) 2NaCl(aq) + Br2(aq)

FClBrI

At

Cl More reactive

Br Less reactive

Solution goesyellow/brown asbromine 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.

FClBrI

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 + ammoniumsulphate 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 reactantsare 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 + sodiumchloride 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 + sodiumchloride 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.

2 Write down the ions in the reactants.



Na+ SO42- Mg2+ NO3

-

Mg2+ SO42- Na+ NO3

-

Both the products are soluble there will be no precipitate.

Sodium + magnesium magnesium + sodiumsulphate 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.

2 Write down the ions in the reactants.



Na+ SO42- Ba2+ NO3

-

Ba2+ SO42- Na+ NO3

-

Barium sulphate is insoluble so there will be a precipitate.

Sodium + barium barium + sodiumsulphate 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-

CO32-

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 a a saltsalt and waterwater.

1 2 14131211109876543

Increasingly acid Increasingly alkali

sodium hydroxidepH 14

hydrochloric acidpH 1

neutralisation

sodium chloridepH 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 sulphatessulphatesNitric acid makes nitratesnitratesHydrochloric 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 + H HCl 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+ + CO32- 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 HH22O O + 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)22 2

Neutralisation equations


Write balanced equations going through the same stages as the previous examples. 1. word equation 2. formulae3. 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 adding Oxidation meant adding oxygen to a substance.oxygen to a substance.

Rusting (iron becoming iron oxide) is an example of oxidation.

Reduction meant taking Reduction 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.

Oxidation is the loss of electrons.

Mg2+

S2-

SMg2+

O2-

O

Mg2+ Cl-Cl-

Cl

MgMg2+

S2-

SMg2+

O2-

O


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

O1½


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

alkali11 reaction of an acid with a

base

• Down1 a solid forms in a solution2 loss of electrons3 competition reaction4 gives solutions containing

H+ ions6 to break down into smaller

particles7 removal of oxygen8 state of balance9 soluble base10 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.

What type of reaction is this?

1. Equilibrium


3. Displacement

4. Precipitation


When the washing soda was added to the lemon juice it fizzed and the pH rose towards 7. What type of reaction is this?

1. Neutralisation


3. Displacement

4. Oxidation


Which of the oxides shown will thermally decompose most easily?

1. Mercury oxide2. Potassium oxide3. Iron oxide4. 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

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© Boardworks Ltd 2003 Chemical Reactions. © Boardworks Ltd 2003 A slide contains teacher’s notes wherever this icon is displayed - To access these notes