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© Boardworks Ltd 2003
Chemical Reactions
© Boardworks Ltd 2003
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© Boardworks Ltd 2003
oxidationand reduction
neutralisation
precipitation reversiblereactions
displacementreactions:
metals
exothermicand endothermic
thermaldecomposition
displacementreactions:
non-metals
Types of chemical change
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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
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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
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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
+
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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 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
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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)
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• 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
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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.
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• 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
Displacement reactions: metals
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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
Displacement reactions: metals
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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
Displacement reactions: metals
iron oxide + aluminium
powder
magnesium fuse
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Here is a photo of the thermit reaction being carried out in a laboratory.
iron oxide + aluminium
powder
magnesium fuse
Displacement reactions: metals
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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
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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.
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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
Displacement reactions: halogens
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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.
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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.
Displacement reactions: halogens
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• 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.
© Boardworks Ltd 2003
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
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• 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
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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) 2NH3 (g) (exothermic)
Low temperature 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) solid–(l) liquid–(g) gas–(aq) aqueous (dissolved in water)
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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
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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
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Another metal hydroxide that precipitates is iron(III) hydroxide. Like many transition metals its compounds are coloured.
Iron + sodium iron + sodiumchloride hydroxide hydroxide chloride
aqueous aqueous solid aqueoussolid
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
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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
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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
aqueous aqueous solid aqueoussolid
Precipitation and solubility
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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.
3 Swap over the + and – ions.
4 Are the products going to be soluble or insoluble?
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
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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.
3 Swap over the + and – ions.
4 Are the products going to be soluble or insoluble?
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
aqueous aqueous solid aqueoussolid
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Separating Precipitates – reminder!
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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
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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
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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
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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
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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 sulphatessulphatesNitric acid makes nitratesnitratesHydrochloric acid makes chlorideschlorides
Eg. Sodium hydroxide + nitric acid forms:
Calcium carbonate + sulphuric acid forms:
Sodium nitrate
calcium sulphate
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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
+
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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
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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
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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
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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
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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
© Boardworks Ltd 2003
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
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• 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
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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.
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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.
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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
© Boardworks Ltd 2003
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
© Boardworks Ltd 2003
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
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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
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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
© Boardworks Ltd 2003
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.
© Boardworks Ltd 2003
• 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
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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
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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
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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
2. Thermal decomposition
3. Displacement
4. 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.
What type of reaction is this?
1. Equilibrium
2. Thermal decomposition
3. Displacement
4. Precipitation
© Boardworks Ltd 2003
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
2. Thermal decomposition
3. Displacement
4. Oxidation
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Which of the oxides shown will thermally decompose most easily?
1. Mercury oxide2. Potassium oxide3. Iron oxide4. Silver oxide
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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
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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 silver nitrate
4.Aluminium and calcium sulphate
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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
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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
© Boardworks Ltd 2003
Which compound can you be sure is soluble in water?
1. Manganese nitrate
2. Osmium iodide
3. Thallium chloride
4. Palladium sulphate