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GCSE Chemistry > Chemical changes > Flashcards

Chemical changes Flashcards

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1
Q

What happens when a metal reacts with oxygen?

A
  • The metal atoms have been oxidised to form a metal oxide.

Eg, Calcium + Copper Oxide -> Calcium oxide + copper

Calcium has been oxidised (gained oxygen)
Copper has been reduced (lost oxygen)

REDUCTION = LOSS OF OXYGEN
OXIDATION = GAIN OF OXYGEN

or

OXIDATION = LOSS OF ELECTRONS
REDUCTION = GAIN OF ELECTRONS

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2
Q

Metal + Water ->

A

Metal Hydroxide (alkali) + Hydrogen

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3
Q

How can we determine the order of reactivity of metals?

A

React them with water and observe how vigorously they react.

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4
Q

State the reactivity series:

A

Potassium
Sodium
Lithium
Calcium
Magnesium
– Carbon (non-metal)
Zinc
Iron
– Hydrogen (non-metal)
Copper

Please Stop Letting Cows Make Zebras In Cages.

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5
Q

Describe the reactivity of Potassium, Sodium and Lithium with water at room temperature.

A

They react rapidly with water at room temperature.

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6
Q

Describe the reactivity of calcium with water at room temperature.

A

Reacts quite rapidly with water at room temperature.

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7
Q

Describe the reactivity of magnesium, zinc, iron and copper with water at room temperature.

A

No reaction.

  • If we want to work out their reactivity, react them with dilute acids if we want to compare their reactivity.
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8
Q

Describe the reactivity of Potassium, Sodium and Lithium with dilute acid.

A

Dangerously fast and vigorous reaction.

  • We should never do this because it is too dangerous because they react too vigorously.
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9
Q

Describe the reactivity of calcium with dilute acid.

A

Extremely vigorous reaction.

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10
Q

Describe the reactivity of magnesium with dilute acid.

A

Rapid reaction.

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11
Q

Describe the reactivity of zinc with dilute acid.

A

Quite rapid reaction.

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12
Q

Describe the reactivity of iron with dilute acid.

A

Quite slow reaction.

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13
Q

Describe the reactivity of copper with dilute acid.

A

No reaction.

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14
Q

What does the reactivity of a metal depend on?

A

Its ability to lose electrons on the outermost shell and form a positive ion.

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15
Q

State why we can find pure gold in the Earth.

A

Because it is so unreactive, so will not react with other metals in the Earth.

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16
Q

True or false, a more reactive metal/element will displace/push out a less reactive element from its compound?

A

True.

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17
Q

How can we extract metals from their ores?

A

Metals less reactive than carbon can be extracted through displacement reactions - using carbon to displace the metal from the metal oxide compound.

We use carbon because other metals are too expensive.

The metal is reduced
The carbon is oxidised

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18
Q

Explain oxidation and reduction in terms of loss and gain of electrons

A

REDUCTION = LOSS OF OXYGEN
OXIDATION = GAIN OF OXYGEN

or

OXIDATION = LOSS OF ELECTRONS
REDUCTION = GAIN OF ELECTRONS

eg S + 2e^- -> S2-
in this case, sulfur has been reduced, because it has gained electrons

eg Mg -> Mg2+ e^-

In this case, magnesium has been oxidised because it has LOST electrons.

OIL
RIG

^above reactions are called half equations. - should also be able to identify oxidation and reduction reactions in symbol equations.

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19
Q

What state symbol are acids followed by?

A

aq (aqueous) because they are dissolved in water.

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20
Q

Ionic equations and displacement reactions practice questions

A
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21
Q

Chemical formula for

Hydrochloric acid
Sulfuric acid
Nitric acid

A

HCl (aq)
H2SO4 (aq)
HNO3 (aq)

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22
Q

Acids react with metals to produce?

A

Salt and hydrogen.

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23
Q

In aqueous solutions, what ions do acids release?

A

a H+ ion.

eg

They ionise and release a H+ ion.

HCL -> H+ + Cl-

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24
Q

What are alkalis and bases?

A

A base is a chemical which can neutralise acids to produce salt and water.

eg Iron(iii)hydroxide

An alkali is an soluble base.
eg Sodium hydroxide

Bases are usually metal hydroxides and metal oxides.

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25
Q

In aqueous solutions, what ions do alkalis/bases release?

A

Hydroxide, OH- ion.

eg NaOH -> Na+ + OH-

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26
Q

How can we determine the acidity or alkalinity of a solution?

A
  • By observing its pH
  • Using a pH probe (electronically)
  • or a universal indicator (chemical which uses colour)
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27
Q

pH scale, what is it?

A

Measure of a solution’s acidity or alkalinity.

0-6 = acid —– red
7 = neutral —– green
8 - 14 = alkaline — purple

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28
Q

True or false, all acids contain hydrogen?

A

True

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29
Q

When acids react with metals, what happens?

A
  • If metal is more reactive than hydrogen in acid, metal will displace hydrogen in acid.
  • This will produce a salt and hydrogen gas
30
Q

Give two examples of acids

A

Hydrochloric acid (produce salts which end in chloride)
Sulfuric acid (produce salts which end in sulfate)

eg sulfuric acid + magnesium -> magnesium sulfate and hydrogen

31
Q

Explain, in terms of oxidation and reduction to explain what happens when acids react with metals.

A
  • Acids release H+ ion
  • When metals react they lose electrons so are oxidised.
  • H+ ion released from acid accept the electrons so H+ ion is reduced.

example:

2HCl + Fe -> FeCl2 + H2

HCl - gives of H+
Fe - gives off 2 electrons

Fe -> Fe2+ + 2e- OXIDATION
2H + 2e- -> H2 REDUCTION

so to explain the rate of reaction with metals and acids:

Magnesium reacts very rapidly bc it easily forms Mg2+ ion
Zinc reacts quite easily bc it easily forms Zn2+ ion
Iron reacts slowly because it less easily forms Fe2+ ion

32
Q

Reactions of acids with alkalis/bases info:

A

Acid + base/alkali -> salt + water

this is a neutralisation reaction.

the salt contains a positive ion from the base/alkali and a negative ion from the acid.

33
Q

Reactions of acids with metal carbonates.

A

Acid + metal carbonate -> salt + water + carbon dioxide.

Examples of metal carbonates

Sodium carbonate
Calcium carbonate
Potassium carbonate

34
Q

Describe a practical to make a soluble salt by using an acid and a metal oxide/carbonate.

A
  • Start with a fixed a fixed volume of dilute sulfuric acid (fixed so acid is limiting reactant and gets completely used up - this is important bc any left over acid would contaminate the salt)
  • Heat acid in a beaker using a Bunsen burner until almost boiling (ensure not boiling otherwise acid will boil over when other reactants are added)
  • Use a spatula to add a small amount of copper oxide to the acid.
  • Gently stir the solution using a glass rod.
  • Copper oxide will react and seem to disappear and the solution will turn blue (the colour of copper sulfate).
  • Continue to add copper oxide into solution if solution remains clear blue and stir.
  • Stop adding if some powder remains after stirring.
  • When powder remains it means reaction is over as all acid has reacted (so solution is now neutral) and copper sulfate has been formed.
  • Filter unreacted copper sulfate using filter funnel and paper.
  • Take copper sulfate solution and place into an evaporating basin.
  • Heat evaporating basing gently over a beaker of boiling water,
  • Heat until approx 1/2 of solution remains
  • Leave solution for 24hrs in cool conditions for copper sulfate crystals to form.
  • Next day, when formed, scrape copper sulfate crystals onto a paper towel and gently pat crystals dry.
35
Q

What is meant by a strong acid?

A

An acid that fully ionises in aqueous solutions to release H+ ions.

36
Q

Examples of strong acids and their ionic equations x3

A

Hydrochloric acid
HCl -> H+ + Cl -

Sulfuric acid
H2SO4 -> 2H+ + SO4-

Nitric acid
HNO3 -> H+ + NO3-

we can tell an acid is strong when arrow is only in one direction.

37
Q

What is meant by a weak acid?

A

Acids that partially ionise in aqueous solutions

  • can be identified from reversible reaction arrow.
38
Q

Examples of weak acids:

A

Carbonic acid
Ethnoic acid
Citric acid

39
Q

What does the pH scale give us an idea of?

A

The concentration of hydrogen ions in a solution.

40
Q

Information about strong and weak acids:

A

Strong acids have a lower pH than weak acids for a given concentration of the acid.
This is because they fully ionise and produce a greater concentration of hydrogen ions than weaker acids which partially ionise.

41
Q

As the pH scale decreases by one unit, the concentration of hydrogen ions increases by how many times?

A

10 times.

eg pH 1 in comparison to pH 2 would have a concentration of hydrogen ions x10 larger than pH 2

each ‘10 times’ is called 1 order of magnitude.

so in comparison to pH 3 it would be 100x more concentrated with hydrogen ions (2 order of magnitude).

10x = 1 order of magnitude
100x = 2 orders of magnitude

etc

42
Q

What does the concentration of an acid tell us?

A

The amount of acid molecules in a given volume of solution.

So a dilute acid will have fewer acid molecules in a given volume than a concentrated acid even if the strength of the acid (how well it ionises) is the same.

So concentration is NOT the same as strength.

43
Q

What are titrations used for?

A

To find the concentration of acid needed to neutralise a certain volume of alkaline solution.

44
Q

Describe a procedure for carrying out a titration.

A
  • Use a pipette and a pipette filler to transfer 25cm^3 of sodium hydroxide into a conical flask
  • When using pipette and pipette filler allow liquid to drain out rather than blowing it out with the filler as this would give an incorrect volume.
  • Add 5 drops methyl orange indicator into the alkali in the conical flask.
  • Place the conical flask on a white tile in order to see the colour change more clearly.
  • Fill burrette with sulfuric acid until acid becomes neutral
  • We can detect this once a colour change begins to occur.
  • When we notice a colour change begin to add acid into alkali drop by drop until colour change from yellow to permanent red.
  • Now read the final volume of acid added from burette
  • when reading off volume ensure eye is levelled with the liquid and ready the burette off from the bottom of the meniscus
  • repeat the titration several times until you get two reading between 0.1cm^3
  • then take the mean value of these for the final volume.
45
Q

To convert from dm^3 to cm^3 :

A

divide by 1000

46
Q

To find conc in g/dm^3 what should we do?

A

multiply the concentration in mol/dm^3 by relative formula mass

47
Q

complete titration calculations

A
48
Q

suggest why solid ionic compounds cannot conduct electricity.

A

Because the ions are locked in place and held together by strong electrostatic forces of attraction so are not free to move and carry charge so can conduct electricity.

49
Q

Suggest why solid ionic compounds can conduct electricity when molten or dissolved in water?

A

Because the strong electrostatic forces of attraction have been broken so ions are free to move and carry charge, so can conduct electricity.

50
Q

What is an electrolyte?

A

A liquid that can conduct electricity.

51
Q

Electrolysis intro, summary.

A

Negative electrode = cathode, which is attached to the positive terminal of the power pack - so ‘covered with electrons’ - this means positive ions are attracted to it and here they are reduced (lose e-) to form an element.

Positive electrode = anode, which is attached to the negative terminal of the power pack, this means negative ions are attracted to it and here they are oxidised (gain e-) to form an element.

So during electrolysis ions are discharged to form an element.

52
Q

What are electrodes made from?

A

Metal or graphite (materials that can conduct electricity)

53
Q

What is meant by ‘inert’

A

refers to a substance or element that is chemically inactive or unreactive under specific conditions.

54
Q

what is electrolysis used for?

A

To extract metals that are more reactive than carbon, from their compounds.

55
Q

Why is aluminium oxide mixed with cryolite?

A

Because aluminium has a v high melting point
cryolite lowers the melting point
this saves energy needed and money.

56
Q

Half equation for Oxygen ion at anode.

A

2O^2- -> O + 4e-

OR 2O^2- - 4e- -> O2

Oxygen ion is loses as it gains 2 electrons (at anode, bc they are attracted to positive anode), so oxygen atom is produced - which pair up.

57
Q

Half equation for Aluminium ion at cathode.

A

Al3+ +3e- -> Al

Aluminium ions are reduced as they gain two electrons, from cathode (which is covered in electrons), so aluminium, atoms are produced.

58
Q

Why must the anode be replace regularly?

A

Electrodes are made from graphite which is a form of carbon and are at 900 deg celc
so oxygen molecules produced at anode reacts with carbon in graphite to produce carbon dioxide.

O2 + C -> CO2

59
Q

State 2 reasons why electrolysis is expensive.

A
  • melting compound is expensive
  • a lot of energy is required to produce electric current.
60
Q

Describe how aluminium is extracted by electrolysis from aluminium oxide.

A
  • Mix molten aluminium oxide with compound called cryolite.
  • Aluminium has a very high melting point, cryolite lowers melting point of aluminium - this reduces energy need and saves money.
  • Set up the cathode, anode and power station like this: draw
  • Electrodes should be made from graphite, which conducts electricity and has a very high melting point so can be used at very high temperatures.
  • Place molten aluminium oxide in set up and turn on power station.
  • Positive aluminium ions are attracted to the negative cathode - here they are reduced and gain three electrons to form an aluminium molecules.
  • Negative oxygen ions are attracted to positive anode, here they are oxidised and lose two electrons to form an oxygen molecules.
61
Q

What can aluminium be used for?

A
  • It has a low density so can be used to make drink cans and can be used in air craft.
62
Q

True or false?

A

When a simple ionic compound (eg lead bromide) is electrolysed in the molten state using inert electrodes, the metal (lead) is produced at the cathode and the non-metal (bromine) is produced at the anode.

63
Q

During electrolysis of aqueous solutions, under what conditions is hydrogen produced at the cathode?

A

If the metal is MORE reactive than hydrogen.

So if metal ion is Cu2+ , copper is LESS reactive than hydrogen, so copper will be produced at the cathode.

64
Q

Why must the electrodes be inert during electrolysis?

A

So they don’t react with the molecules produced.

65
Q

Electrolysis of aqueous solution information:

A

An aqueous solution is a solution dissolved in water.
Water ionises to produce a H+ and OH- (hydrogen and hydroxide) ion - this reaction is reversible.
So if we have aqueous copper sulfate solution; CuSO4 - ions from this will be Cu2+ and SO4^2-
So during the entire process of this electrolysis the ions to be considered are:
- H+
- OH-
- Cu2+
- SO4^2-

the question? which positive ion will be reduced at the cathode?

RULE; if metal ion is MORE reactive than hydrogen, hydrogen will be produced at the cathode (hydrogen ion will be reduced)

So copper is LESS reactive than hydrogen, so copper ion will be reduced at cathode; copper will be produced.

AT ANODE; oxygen is always produced during electrolysis of aqueous solutions, except for when aqueous solutions contains halide ions, eg chloride.

HALF EQUATIONS:

copper is produced at cathode:
reduction
Cu^2+ -> 2e- ————-> Cu

oxygen is produced at the anode
oxidation
4OH- ————–> O2 + 2H2O + 4e-

OR
oxidation
4OH- - 4e- ———-> O2 + 2H2O

66
Q

Half equation when copper is produced at the cathode during electrolysis of CuSO4:

A

copper is produced at cathode:
reduction
Cu^2+ -> 2e- ————-> Cu

67
Q

Half equation when oxygen is produced at anode during electrolysis of CuSO4:

A

oxidation
4OH- ————–> O2 + 2H2O + 4e-

OR
oxidation
4OH- - 4e- ———-> O2 + 2H2O

68
Q

Electrolysis of aqueous solutions containing a halogen, eg Sodium chloride (NaCl)

A

Ions to consider:
- Na+
- Cl-
- H+
- OH-

what will be produced at the cathode?
Hydrogen, because sodium is more reactive than it.

What will be produced at the anode?
Halogen molecule, so in this case, chlorine because halogen ion is in aqueous solution.

HALF EQUATIONS

At cathode
reduction
2H+ +2e- ————-> H2

At anode
oxidation
2Cl - —————-> Cl2 + 2e-

OR oxidation
2Cl- -2e- —————->Cl2

69
Q

Half equation at cathode of electrolysis of NaCl

A

At cathode
reduction
2H+ +2e- ————-> H2

70
Q

Half equation at anode of electrolysis

A

At anode
oxidation
2Cl - —————-> Cl2 + 2e-

OR oxidation
2Cl- -2e- —————->Cl2

71
Q

Oxygen is only produced at anode when?

A

Aqueous solution doesn’t contain halogen. If it does then the halide ion will be oxidised to produce a halogen at the anode.

72
Q

Describe how to investigate what happens when aqueous solutions undergo electrolysis.

A
  • Pour 50cm^3 of copper chloride solution into a beaker.
  • Place a plastic Petri dish with two holes over the beaker
  • Insert a carbon graphite rode (electrodes) into each hole
  • These electrodes are inert and must not touch each other as this would create a short circuit.
  • Add crocodile leads to the rods and connect the rods to a low voltage power supply and switch this on.
  • Observe the electrodes.
  • We should see at the anode bubbles from a gas will be produced, this gas is chlorine
  • Chlorine gas is produced because halide ion will always be discharged during electrolysis when aqueous solution contains a halogen.
  • We can confirm whether this gas is chlorine by testing it with damp blue litmus paper, if it bleaches then gas is chlorine.
  • Repeat this process using another of the same set up, using 50cm ^3 of sodium chloride solution
  • Observe the electrodes
  • We should see that at the anode a gas is produced, this is noticed from bubbles
  • This gas is chlorine so bleaches litmus paper
  • At cathode, we will also notice gas bubbles, this means hydrogen is produced, hydrogen is produced at the cathode because sodium is more reactive than hydrogen
  • To prove this gas is hydrogen we can conduct the squeaky pop test (collect gas in a test tube and place a lit splint near the opening of test tube containing gas, if you hear a squeaky pop sound, H2 is present)

GCSE Chemistry (13 decks)

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