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Z AQA GCSE chemistry flash cards > topic 4 Chemical Changes > Flashcards

topic 4 Chemical Changes Flashcards

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1
Q

what ion is produced by all alkalis in aqueous solutions?

A

OH- Hydroxide ions

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2
Q

B adds some citric acid to some tap water. What happens to the concentration of the H+ ions in the water? What happens to the pH of the water?

A

the conc. of H+ ions inc. because acids form H+ ions in water. The pH decreases as the water becomes acidic.

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3
Q

Why are wide range indicators (like universal indicator) suitable for estimating the pH of a solution?

A
  • change colour gradully over a broad pH range
  • each colour corresponds to a different pH
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4
Q

Other than using universal indicator, how else can you determine the pH of a solution?

A

pH probe/meter

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5
Q

Which type of reaction occurs when hydrochloric acid reacts with potassium hydroxide?

A
  • a neutralisation reaction
  • acid-base neutralisation reactions always prodce a salt and water
  • in this reaction potassium chloride (KCl) and water are produced
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6
Q

T or F? when carrying out a titration you must always add the acid to the alkali.

A

False you can also add the alkali to the acid

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7
Q

what is meant by the end point of a titration?

A

the point at which the indicator changes colour - so the point at which the acid (or alkali) has been fully neutralised

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8
Q
  1. S needs to accurately measure out 50cm3 of an alkali. He has a 50cm3 pipette and a 50cm3 measuring cylinder. Which should he use to measure out the alkali?
A

The 50cm3 pipette. Pipettes are calibrated, which reduces transfer errors. This makes them more accurate than measuring cylinders.

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9
Q

Explain how you could use a burette to measure the volume of an acid required to neutralise an alkali.

A

fill the burette with acid, and measure the initial volume. Add the acid to the alkali until the end-point is reached, and measure the final volume. Then take the initial reading away from the final reading to calculate the volume of of acid you added.

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10
Q

Why should you always pour acid into a burette below eye level?

A

there’s a chance you may spill some when filling the burette, so having your eyes above the level of the acid will help prevent it from getting in your eyes

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11
Q

Why should you repeat a titration until you obtain several consistent results?

A

to inc. the accuracy and precision of your titration and identify any anomalous results

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12
Q

Luke is carrying out a titration to calculate the concentration of an alkali. He plans to use universal indicator to determine the end-point. Why is this a bad idea?

A

Universal indicator gradually changes colour as the pH of a solution changes. In a titration you want to see a sudden, sharp colour change at the end-point.

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13
Q

4) Reeta carries out a titration four times. The volumes of acid she adds are: 23.3 cm³, 21.8 cm³, 23.4 cm³, and 23.3 cm³. Calculate the mean volume of acid she added, ignoring any anomalous results.

A

4) Mean volume = (23.3 + 23.3 + 23.4) 3 = 70 ÷ 3 = 23.3 cm³ (to 3 significant figures) In a titration, you ideally want to measure volumes that are within 0.10 cm³ of each other. 21.8 cm³ is anomalously low compared to the other values, so you should exclude it from your calculation.

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14
Q

5) The equation for the reaction between sulfuric acid and sodium hydroxide is:
2NaOH + H₂SO4 → Na₂SO4 + 2H₂O
25 cm³ of 2 mol/dm³ sodium hydroxide solution is neutralised in a titration by
50 cm³ of sulfuric acid. Calculate the concentration, in mol/dm³, of sulfuric acid used in the titration.

A

5) Moles of NaOH = concentration x volume
= 2 x (25 ÷ 1000) = 2 × 0.025 = 0.05 mol

Ratio of NaOH to H₂SO4 is 2:1, so Moles of H₂SO4 = 0.05 ÷ 2 = 0.025 mol Concentration of H₂SO, = moles/volume = 0.025 + (50÷1000) = 0.025/0.05
= 0.5 mol/dm³

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15
Q

Give some examples of strong and weak acids

A

Strong acids: hydrochloric acid, sulphuric acid and nitric acid
Weak acid: ethanoic acid, citric acid and carbonic acid.

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16
Q

T or F? All weak acids have low concentrations.

A

False. You can have both high and low concentrations of weak acids. Acid concentration is the number of acid molecules dissolved per unit volume. Acid strength is the proportion of those molecules that have dissociated to release H+ ions.

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17
Q

Explain what is meant by a weak acid.

A

A weak acid is an acid that partially ionises in a solution. Only a small proportion of the acid particles dissociate to release H+ ions.

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18
Q

A student reacts a weak acid and a strong acid with identical strips of magnesium. The reaction produces a gas. If both the acids have the same volume and concentration, which reaction will produce 50cm3 of gas in the shortest amount of time. Explain your answer.

A

The reaction with the strong acid. For a given concentration, the stronger the acid, the more reactive it is, so the faster it will react. This is because the reactions of acids involve H+ ions, and strong acids release more H+ ions than weak acids do. So in strong acids the concentration of H+ ions is higher than in weak acids.

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19
Q

S adds some acid to a solution. The pH of the solution decreases by 2. How many times greater is the concentration of H+ ions in the solution, now that S has added some acid to it?

A

For every decrease in 1 on the pH scale, the increase in H+ ions increases x10. So a decrease in 2 means that the H+ ion concentration is 10 x 10 = 100 times greater.

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20
Q

T or F? Bases that are insoluble in water do not take part in neutralisation reactions with acids.

A

False. Insoluble bases react in neutralisation reactions if added to acids.

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21
Q

What are the products when an acid reacts with a metal hydroxide?

A

A salt + water (metal oxides and hydroxides are bases so you get a neutralisation reaction).

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22
Q

What are the products when an acid reacts with a metal hydroxide?

A

A salt + water (metal oxides and hydroxides are bases so you get a neutralisation reaction).

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23
Q

Name the products of the reaction between sulphuric acid and copper oxide?

A

Copper sulphate and water

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24
Q

M wants to make a sample of potassium nitrate. Suggest the reagents she could use to produce this sample.

A

potassium hydroxide and nitric acid

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25
Q

What is the chemical formula of the salt that forms when hydrochloric acid (HCl) reacts with calcium carbonate (CaCO3). What are the other products of this reaction?

A

CaCl2 (calcium chloride). The other products are water (H2O) and carbon dioxide (CO2)

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26
Q

Briefly describe how you could make a pure dry sample of copper chloride (a soluble salt) from hydrochloric acid and copper oxide (an insoluble base).

A

e.g. warm the acid gently using a Bunsen burner or water bath, then turn off the heat source. Add the copper oxide until no more reacts, making sure you stir it. Filter out the excess solid using a filter paper and a funnel. Crystallise the copper chloride salt solution by heating it gently in a water bath or electric heater until it becomes more concentrated, then leave to cool. Crystals of copper chloride will form. Filter these out of the solution and leave to dry.

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27
Q

which two non-metals are often included in the reactivity series?

A

carbon and hydrogen

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28
Q

Put lithium, iron and sodium in increasing order of their reactivity with water.

A

Increasing order of reactivity: iron, lithium, sodium

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29
Q

true or false. Elements at the top of the reactivity series form positive ions more easily than those at the bottom.

A

true. The metals at the top of the reactivity series (like potassium and sodium) are ones that can easily lose electrons, so they form stable positive ions more easily.

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30
Q

What are the products of the reaction between sodium and water?

A

Sodium hydroxide and (NaOH) and Hydrogen (H2)

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31
Q

A student adds samples of copper, magnesium, zinc and iron into individual samples of warm hydrochloric acid and measures the temperature change of the acid in each case. Her results are: copper (4°C), magnesium (58°C), zinc (10°C). Predict the temperature change she observed for the sample of iron. Explain your answer.

A

Any value between 4°C and 10°C. The more reactive the metal, the greater the increase in temperature of the acid. iron is above copper in the reactivity series but below zinc. So iron will cause a larger temperature change than copper but a smaller one than zinc.

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32
Q

D adds some magnesium to a test tube containing dilute hydrochloric acid. he then adds some iron to another test tube containing an identical sample of the acid. Predict what D will observe in both cases. Explain your answer.

A

The reaction with the magnesium will be vigorous producing lots of bubbles. The reaction with iron will be slower producing fewer bubbles. The reaction is more vigorous with magnesium because magnesium is above iron in the reactivity series. This also means that magnesium reacts faster, so forms more bubbles in the same amount of time.

33
Q

Give an example of a metal that can be mined in its elemental form.

A

gold - it’s so unreactive it doesn’t react with oxygen.

34
Q

How is iron extracted from iron oxide?

A

Iron is extracted by reduction with carbon. Iron can be extracted this way because it is below carbon in the reactivity series.

35
Q

Explain why some metals don’t need to be extracted from ores.

A

some metals are found in the earth as just the pure metal, because they are so unreactive.

36
Q

Explain why magnesium can’t be extracted from its ore by reduction with carbon.

A

Carbon can only reduce metals that are below it in the reactivity series, and magnesium is more reactive than carbon. (It has to be extracted by carbon instead)

37
Q

Zinc can be extracted from its ore by either electrolysis or reaction with carbon. Suggest why a mining company might choose to extract zinc using reduction with carbon rather than electrolysis.

A

electrolysis is an expensive process

38
Q

6) Which substances are oxidised in each of the following reactions?
A. CuO + H₂ → Cu + H₂O

B. 2Mg + O₂ → 2MgO
Explain your reasoning in each case.

A

6)

A. H₂/hydrogen is oxidised, because it gains oxygen.

B. Mg/magnesium is oxidised, because it gains oxygen.

39
Q

What is the chemical formula of the product formed from the oxidation of calcium?

A

CaO. The calcium is oxidised, so it gains oxygen to form calcium oxide.

40
Q

What sort of reaction occurs when zinc is added to a solution of copper chloride?

A

a displacement/ redox reaction

41
Q

What compound is formed when magnesium is added to solution of zinc sulphate?

A

Magnesium sulphate (MgSO4). This is because magnesium is more reactive than zinc, so will displace it from zinc sulphate.

42
Q

The table shows whether a reaction occurs when different metals were added to solutions of metal nitrates. Explain what the results show.

A

More reactive metals can displace less reactive metals from their compounds. Calcium is more reactive than both iron and copper, so can displace them from iron nitrate and copper nitrate respectively. Iron is more reactive than copper so can displace copper from copper nitrate. Iron is less reactive than calcium so no reactions occur between iron and calcium nitrate. Copper is less reactive than both calcium and iron , so no reactions occur when copper is added to either copper nitrate or calcium nitrate.

43
Q

Explain, in terms of electrons, what happens to a species that is reduced in a chemical reaction.

A

A species that is reduced gains electrons.

44
Q

5) The ionic equation for the reaction between magnesium and iron(III) chloride is:
3Mg(s)+ 2Fe³+(aq)→3Mg2+(aq) +2Fe(s)

Which species is reduced and which species is oxidised in this reaction?
Explain your answer.

A

5) Fe³+ is reduced, as it goes from a positive ion to a neutral atom, showing that it has gained electrons. Mg is oxidised, as it goes from a neutral atom to a positive ion, showing that it has lost electrons.

45
Q

What are half equations?

A

Half equations show you what happens at one of the electrodes during electrolysis. They are also used in redox reactions. They involve either atoms or molecules gaining or losing electrons to become ions, or ions gaining or losing electrons to become atoms and molecules.

46
Q

How do write a half equation? e.g. conversion of oxide ions into an oxygen molecule.

A

Oxide ion to oxygen molecule:

i) Write down the reactant and the product: O2- ⇢ O2
ii) Balance the atoms: 2O2- ⇢ O2
iii) write the total charge under each species in the equation:

2O2- ⇢ O2

(4e-) (0)

iv) Balance the charge by adding electrons: 2O2- ⇢ O2 + 4e-

47
Q

What do you need to remember when balancing a half equation?

A

You should have an equal charge and the same number of atoms on both sides.

48
Q

What happens during anode reactions?

What happens during cathode reactions?

A

During anode reactions (the loss of electrons) electrons are added to the side with the ion.

During cathode reactions (the gain of electrons) electrons are added to the side with the neutral molecule.

49
Q

Write the balanced half equations for the following anode reactions:

  • chlorine
  • oxygen
  • oxygen form hydroxide ions Write the half equations for the following cathode reactions:
  • sodium
  • lead
  • hydrogen
A

Chlorine: 2Cl- → Cl2 + 2e-

Oxygen: 2O2- → O2 + 4e-

Oxygen from hydroxide ions): 4OH- → 2H2O + O2 + 4e-

Sodium: Na+ + e- → Na

Lead: Pb2+ + 2e- → Pb

Hydrogen: 2H+ + 2e- → H2

50
Q

What processes take place at the:

a) cathode (-ive)
b) anode (+ive)

A

a) reduCtion takes place at the Cathode
b) oxidAtion takes place at the Anode.

51
Q

Write the half equations for what forms at the electrodes during the electrolysis of aluminium oxide.

A

At the -ive electrode:

  • metals form +ive ions so they’re attracted to the negative electrode.
  • so aluminium is produced at the -ive electrode
  • Al3+ + 3e- → Al At the +ive electrode:
  • non-metals form -ive ions, so they’re attreacted to the +ive electrode
  • so oxygen is produced at the +ive electrode 2O2- →O2 + 4e- Overall equation:
  • AlO3 → Al + O2
  • 2Al2O3 (l) → 4Al (l) + 3O2 (g)
52
Q

1) What name is given to the negative electrode in electrolysis?

A

1) the cathode

53
Q

2) Why is aluminium oxide mixed with cryolite during the extraction of aluminium?

A

2) To reduce the melting point of the aluminium oxide.

54
Q

3) What is an electrolyte?

A

3) A liquid or solution, usually a molten or dissolved ionic compound, that can
conduct electricity.

55
Q

4) Which electrode do the lead ions move towards during electrolysis of molten PbBr₂?

A

4) The cathode/negative electrode. The lead ions (Pb2+) are positively charged,
so they’re attracted to the negative electrode/cathode.

56
Q

5) A student is carrying out an electrolysis of a molten metal oxide using graphite electrodes. After several repetitions of the experiment, she notices that the anode has
begun to degrade. Suggest a reason for this.

A

5) E.g. the anode is reacting with oxygen to produce carbon dioxide, wearing it away.

57
Q

6) What are the products that will be produced at the anode and the cathode in the electrolysis of molten CaCl₂?

A

6) Anode: Cl₂/chlorine
Cathode: Ca/calcium

58
Q

7) Describe the reaction that occurs at the anode during the electrolysis of
molten potassium bromide (KBr). Is this an example of oxidation or reduction?

A

7) The negative Br ions are attracted to the positively-charged anode, where they lose electrons to form elemental bromine, Br₂. Because electrons are lost, this is an example
an oxidation reaction.

59
Q

Things to remember about half equations?

A
  • electons are shown as e-
  • the number of atoms of each element must be the same on both sides
  • the total charge on each side must be the same (usually zero)
60
Q

How does electrolysis of aqueous solutions differ to normal electrolyis?

A

In aqueous solutions you have to consider there are ions in the water.

As well as the ions from the ionic compound there are hydrogen ions (H+) and hydroxide ions (OH-) in the water.

61
Q

What is the ionic equation for water?

A

H2O (l) ⇌ H+ (aq) + OH- (aq)

62
Q

How do the ions in the water affect eletrolysis?

A

Ions in the water may change what forms at each electrode. Which ions will be discharged at the electrodes when the solution is electrolysed will depend on the relative reactivity of all the ions in the solution.

63
Q

What is the order of the reactvity series?

A

Please Send Lions Cats Zebras And Cute Zebras Into Hot Countries Signed Gordon

K Potassium

Na Sodium

Li Lithium

Ca Calcium

Mg Magenesium

Al Aluminium

C Carbon

Zn Zinc

Fe Iron

H Hydrogen

Cu Copper

Ag Silver

Au Gold

64
Q

How do ions in the water affeact what is produced at the cathode (-ive)?

A
  • At the cathode, if H+ ions and metal ions are present, hydrogen gas will be produced if the ions form an elemental metal that is more reactive than hydrogen. (e.g. sodium ions).
  • If the metal ions form an elemental metal that is less reactive than hydrogen (e.g. copper ions) a layer of ther pure metal will be produced instead, which will coat the cathode.
65
Q

How do ions in the water affect what is produced at the anode (+ive)?

A
  • At the anode, if OH- and halide ions (Cl-, Br- and I-) are present, molecules of bromine, chlorine or iodine will be formed.
  • If no halide ions are present, then OH- ions from the water will be discharged as oxygen gas (and water) will be produced.
66
Q

1) Give the names of the ions that are produced by the breakdown of
water molecules.

A

1) hydrogen/H+ and hydroxide/OH- ions

67
Q

2) What can be used as inert electrodes in electrolysis experiments?

A

2) E.g. graphite

68
Q

3) What is the product at the cathode in the electrolysis of aqueous magnesium sulfate? Explain your answer.

A

3) Hydrogen gas/H₂ will be produced, as magnesium is more reactive than hydrogen.

69
Q

4) A student carries out an electrolysis of an ionic compound. The product produced at the anode is chlorine gas. The product produced at the cathode is copper metal. Suggest an identity for the aqueous compound that the student electrolysed.

A

4) copper chloride/CuCl₂

70
Q

5) Give the product at each of the electrodes in the electrolysis of aqueous
sodium iodide. What is the half equation for the oxidation process?

A

5) Product at anode/positive electrode: iodine/I₂
Product at cathode/negative electrode: hydrogen/H2 (Na is more reactive than H)

Half equation for the oxidation process: 2I-→ I2 +2e-

71
Q

6) What are the half equations for the reactions occurring at each electrode in the electrolysis of aqueous copper sulfate (CuSO4)?

A

6) Anode/positive electrode:

40H- → O₂ + 2H₂O + 4e-

Cathode/negative electrode:

Cu²+ + 2e- → Cu

72
Q

1) Why do ionic solids have to be melted before they can be electrolysed?

A

1) In an ionic solid, the ions are in fixed positions and can’t move. In the molten state, the ions are able to move about and can conduct electricity.

73
Q

2) What is the product at the cathode in the electrolysis of molten lead bromide?

A

2) Lead. The positive lead ions move to the cathode and form lead metal.

74
Q

3) Tungsten is below carbon in the reactivity series. Suggest why electrolysis might be used to extract tungsten from its ore, rather than reduction with carbon.

A

3) E.g. tungsten might react with the carbon.

75
Q

5) What is the half equation for the reduction of hydrogen ions?

A

5) 2H+ + 2e- → H₂

76
Q

4) Jessica wants to extract sodium metal from a solid sample of sodium chloride. She dissolves the sample in water and electrolyses it using inert electrodes. Why will this not produce sodium metal? Suggest how the experiment could be altered to obtain a sample of sodium metal.

A

4) In solution, there are both sodium and hydrogen ions present. Because sodium is more reactive than hydrogen, hydrogen gas will be discharged at the cathode and the sodium will remain in the solution as sodium ions. In order to extract the sodium there can’t be hydrogen ions present, so molten sodium chloride should be used instead.

77
Q

6) What is produced by the oxidation process in the electrolysis of aqueous zinc sulfate?
Explain your answer.

A

6) Oxygen (and water) will be produced, because there aren’t any halide ions present.

78
Q

What process takes place at the cathode?

A

Reduction: +ive metal ions gain electrons at the cathode (-ive) to form neutral atoms.

e.g. Pb2+ + 2e- ⇒ Pb

79
Q

What process takes place at the anode?

A

Oxidation: -ive non-metal ions lose electrons at the anode (+ive) to form neutral atoms or molecules.

e.g. 2Br- → Br2

Z AQA GCSE chemistry flash cards (10 decks)

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