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A level Chemistry AQA Lewis > 3.2.3 - Group 7, The Halogens > Flashcards

3.2.3 - Group 7, The Halogens Flashcards

1
Q

What colour is fluorine?

A

Pale yellow

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2
Q

What colour is chlorine?

A

green

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3
Q

What colour is bromine?

A

red-brown

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4
Q

What colour is iodine?

A

grey

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5
Q

What physical state is fluorine at room temperature?

A

gas

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6
Q

What physical state is chlorine at room temperature?

A

gas

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7
Q

What physical state is bromine at room temperature?

A

liquid

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8
Q

What physical state is iodine at room temperature?

A

solid

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9
Q

Boiling points _____ down the group

A

increase

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10
Q

Why do boiling points increase down the group? (3)

A
  • ∵ Van der Waals forces between molecules
  • Increase with size or Mr or surface area
  • More energy needed to overcome these forces

Trend is shown in changes of physical state from fluorine (gas) to iodine (solid))

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11
Q

Electronegativity _______ down the group

A

decreases

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12
Q

Why does electronegativity decrease down the group? (2)

A
  • Shielding increases or the atomic size increases
  • Weaker attraction by nucleus for bonding pair of electrons in the covalent bond
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13
Q

Halogens become _____ oxidising down the group

A

less

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14
Q

Why do halogens become less oxidising down the group?

A
  • Get less reactive down the group ∵ atoms become larger
  • Outer shell further away from nucleus ∴ electrons less attracted to it
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15
Q

How can you see the relative oxidising strengths of halogens?

A

By displacement reactions with halide ions

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16
Q

What is the basic rule for halogens in a displacement reaction with halide ions

A

A halogen will displace a halide from solution if the halide is below it in the periodic table

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17
Q

What can we use these displacement reactions (halogens and halides) for?

A

To help identify which halogen (or halide) is present in solution

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18
Q

Why can’t you investigate fluorine in aqueous solution?

A

∵ it reacts with water

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19
Q

Name 2 uses of chlorine

A
  • Used to kill bacteria in water
  • Mix it with sodium hydroxide to make bleach
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20
Q

Describe how you get bleach

A

If you mix chlorine gas with cold, dilute, aqueous sodium hydroxide, you get sodium chlorate(I) solution (NaClO(aq)) = bleach

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21
Q

Write an equation to show how bleach is made

A
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22
Q

Write the oxidation states of chlorine in this equation

A
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23
Q

What do the oxidation states of chlorine in this equation tell us?

A

Oxidation state of Cl goes up and down, meaning chlorine is both oxidised and reduced. This is called disproportionation.

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24
Q

Name 3 uses of sodium chlorate(I) solution (bleach)

A
  • Water treatment (kills bacteria)
  • Bleach paper and textiles
  • Cleaning toilets
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25
Q

When you mix chlorine with water, what does it undergo?

A

Disproportionation

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26
Q

Write an equation to show the reaction of chlorine with water to form chloride ions and chlorate(I) ions.

Include the oxidation states for chlorine.

A
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27
Q

In sunlight, what can chlorine do to water & what does this form?

A

Can decompose water to from chloride ions and oxygen

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28
Q

Write an equation to show the reaction of chlorine with water to form chloride ions and oxygen

29
Q

______ ____ kill bacteria

A

Chlorate(I) ions

30
Q

Chlorine is ___

31
Q

Name 3 advantages of adding chlorine to water supplies

A

Chlorine is important part of water treatment

  • Kills pathogens
  • Some chlorine persists in water and prevents reinfection further down supply
  • Prevents growth of algae, eliminating bad tastes and smells, & removes discolouration caused by organic compounds
32
Q

Describe disadvantages of adding chlorine to water supplies (3)

A

Risks from using chlorine to treat water

  • Wasteful as most potable water not used for drinking- used in washing clothes
  • Some people suffer eye irritation
  • Can react with organic compounds to produce harmful substances (carcinogenic)
33
Q

Reducing Power of Halides ______ Down the Group

34
Q

Why does the reducing power of halides increase down the group?

A

(To reduce something, halide ion needs lose electron from its outer shell)

  • As you go down group, attraction between outer electron gets weaker ∵
    • Larger atomic radius
      • So electrons further away from +ve nucleus
    • More shielding
  • So further down group halide ion is = easier it loses electrons & greater its reducing power
35
Q

Describe the test for halides

A
  1. Add dilute nitric acid (HNO3) and then silver nitrate solution (AgNO3)
  2. Precipitate is formed (of silver halide)
36
Q

Test for Halides

What is the result for fluoride?

A

No precipitate

37
Q

Test for Halides

What is the result for chloride?

A

White precipitate of silver chloride

38
Q

Test for Halides

What is the result for bromide?

A

Cream precipitate of silver bromide

39
Q

Test for Halides

What is the result for iodide?

A

Yellow precipitate of silver iodide

40
Q

Test for Halides

Why do you add dilute nitric acid?

A

To remove ions which may interfere with test

41
Q

Test for Halides

Write the general equation

42
Q

Fill in the gaps (speed)

43
Q

Test for Halides

What can you do to be sure of your results & why does this work?

A
  • You can test your results by adding ammonia solution
  • Each silver halide has a different solubility in ammonia
44
Q

Fill in the gaps

45
Q

What reactions reflect the trend in the reducing ability of halide ions?

A

Reactions of solid sodium halides with concentrated sulfuric acid

46
Q

All halides react with concentrated sulfuric acid to give what as a product?

A

hydrogen halide

47
Q

Write the equations for the reaction of NaF or NaCl with H2SO4

48
Q

Reaction of NaF or NaCl with H2SO4

Describe the observations

A

Misty (white) fumes hydrogen fluoride / hydrogen chlroide

49
Q

Name 2 features of these reactions

A
  • HF and HCl aren’t strong enough reducing agents to reduce sulfuric acid so reaction stops there
  • Acid –base reactions & not redox reactions
    • H2SO4 plays the role of an acid
    • Oxidation states of halide and sulfur stay the same (-1 and +6)
50
Q

Write the equations for the reaction of NaBr with H2SO4

A
  1. Acid-base step
  2. Redox step
51
Q

Describe the observations in this reaction

A

Misty fumes of hydrogen bromide gas (HBr)

52
Q

Write the oxidation states of S & Br for the 2nd equation

53
Q

Describe the reducing agent in these reactions

A

HBr is stronger reducing agent than HCl and reacts with H2SO4 in a redox reaction

54
Q

Describe the observations in the 2nd equation

A

Choking fumes of SO2 and orange fumes of Br2

55
Q

Write the equations for the reaction of Nal with H2SO4

56
Q

Reactions of NaI with H2SO4

Write the oxidation state of S & I

57
Q

Reactions of NaI with H2SO4

Write the oxidation state of S & I

58
Q

Describe the reducing agent in these equations

59
Q

Bromine reacts with phosphorus to form phosphorus tribromide. Write an equation for this reaction.

A

6Br2 + P4 → 4PBr3

60
Q

Describe the role of H2SO4 in these reactions

A
  1. H2SO4 plays the role of acid in the first step
  2. Then acts as an oxidising agent in the second redox
61
Q

State the reduction product

A

Sulfur dioxide

62
Q

State the oxidation and reduction half equations for when bromine reacts with conc. sulfuric acid

A
  • Ox ½ equation: 2Br- → Br2 + 2e-
  • Re ½ equation: H2SO4 + 2H+ + 2 e- → SO2 + 2H2O
63
Q

Reactions of NaI with H2SO4

Write an equation that gives sulfur

A

6HI + H2SO4 → 3I2 + S + 4H2O

64
Q

Reactions of NaI with H2SO4

Name the products

A

Sulfur dioxide, sulfur and hydrogen sulfide

65
Q

Reactions of NaI with H2SO4

Describe H2SO4 role in the reactions

A

H2SO4 plays the role of acid in the first step producing HI & then acts as an oxidising agent in the three redox steps

66
Q

Reactions of NaI with H2SO4

Write the oxidation half equation and 3 reduction half equations.

A

Oxidation

  • 2I- → I2 + 2e-

Reduction:

  • H2SO4 + 2H+ + 2e- → SO2 + 2H2O
  • H2SO4 + 6H+ + 6e- → S + 4H2O
  • H2SO4 + 8H+ + 8e- → H2S + 4H2O
67
Q

Reactions of NaI with H2SO4

Describe the observations

A
  • Sulfur dioxide is a choking gas
  • Sulfur is a yellow solid
  • Hydrogen sulfide has a smell of bad eggs
68
Q

2NaBr + 2H2SO4 → Na2SO4 + Br2 + SO2 + 2H2O

Explain why bromide ions reacts differently from chloride ions (2)

A
  • Br− ions are bigger than Cl− ions
  • ∴ Br− ions are more easily oxidised/lose an electron
69
Q

Write an ionic equation for the reaction between chlorine and cold dilute sodium hydroxide solution

A

Cl2 + 2OH- → ClO- + Cl- + H2O

A level Chemistry AQA Lewis (34 decks)

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