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3.2 A level Inorganic Chemistry > 3.2.3 Group 7 (17) The Halogens > Flashcards

3.2.3 Group 7 (17) The Halogens Flashcards

1
Q

What is a halide?

A

The negative ion of a halogen
X-

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2
Q

How does boiling point change descending group 7?

A

Boiling points increase going down the group.

This trend is shown by the changes of physical state from fluorine to iodine

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3
Q

Why does boiling point increase going down group 7

A

This is due to the increasing strength of van der waals forces as the size and relative mass of the molecules increases.

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4
Q

How does electronegativity change descending group 7?

A

Electronegativity decreases down the group

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5
Q

Why does electronegativity decrease down group 7?

A

Electronegativity is the tendency of an atom to attract a bonding pair of electrons. The halogens are all highly electronegative elements. But larger atoms attract electrons attract electrons less than smaller atoms. This is because the electrons are further from the nucleus and are shielded by more electrons.

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6
Q

Why do halogens get less reactive going down group 7?

A

When halogens react they gain an electron.
They get less reactive down the group, because the atoms become larger. The outer shell is further from the nucleus, so electrons are less strongly attracted to it.
The halogens become less oxidising descending the group

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7
Q

How can we observe the relative oxidising strength of halogens?

A

In displacement reactions, a halogen will displace a halide from solution if the halide is bellow it in the periodic table.
If a halide is displaced by a halogen then it was less oxidising than the halogen.

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8
Q

Why are displacement reations useful?

A

They can be used to determine the identity of a halide is present in a solution

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9
Q

What happens when you add chlorine water (colourless) to potassium chloride (colourless)?

A

No reaction

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10
Q

What happens when you add chlorine water (colourless) to potasium bromide (colourless)?

A

Orange solution (Br2 formed)

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11
Q

What happens when you add chlorine water (colourless) to potassium iodide solution (colourless)?

A

Brown solution (I2) formed

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12
Q

What happens when you add bromine water (orange) to potassium chloride (colourless)?

A

No reaction

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13
Q

What happens when you add bromine water (orange) to potassium bromide (colourless)

A

No reaction

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14
Q

What happens when you add bromine water (orange) to potassium ioddie (colourless)

A

Brown solution (I2) formed

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15
Q

What happens when you add iodine solution (brown) to potassium chloride (colourless)?

A

No reaction

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16
Q

What happens when you add iodine solution (brown) to potassium bromide (colourless)?

A

No reaction

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17
Q

What happens when you add iodine solution (brown) to potassium iodide (colourless)?

A

No reaction

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18
Q

How do you make bleach

A

By mixing chlorine gas with cold, diute, aqueous sodium hydroxide sodium chlorate(I) solution, NaClO(aq), aka bleach is formed

2NaOH(aq) + Cl2(g) -> NaClO(aq) + NaCl(aq) + H2O(l)

In this reaction, the oxidation state of chlorine goes up and down, therefore chlorine is both oxidised and reduced. This is called disproportionation

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19
Q

Name some uses for bleach

A
  • Water treatment
  • Bleaching paper and textiles
  • Cleaning toilets
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20
Q

How does beach kill bacteria?

A

When chlorine is mixed with water it undergoes disportionation, you get a mixture of chloride and chlorate(I) ions:

Cl2(g) + H2O(l) <=> 2H+(aq) + Cl-(aq) + ClO-(aq)

In sunlight, chlorine can also decompose water to form chloride ions and water:

2Cl2(g) + 2H2O(l) <=> 4H+(aq) + 4Cl-(aq) + O2(g)

Chlorate(I) ions kill bacteria

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21
Q

Give the advantages of using chlorine to treat water

A
  • It kills disease-causing microorganisms
  • Some chlorine persists in the water and prevents reinfection further down the supply
  • It prevents growth of algae, eliminating bad tastes and smells, removes discolouration caused by organic compounds
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22
Q

Give the risks of using chlorine to treat water

A
  • Chlorine gas is very harmful if it’s breathed in - it irritates the respiritory system. Liquid chlorine on the skin or eyes cn cause severe chemical burns. Accidents involving chlorine can be serious and fatal
  • Water contains a variety of organic compounds (e.g. from the decomposition of plants). Chlorine reacts with these compounds to form chlorinated hydrocarbons. Many chlorinated hydrocarbons are carcinogenic.
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23
Q

Why is chlorine used to treat water despite the risks?

A

The risk of cancer is much smaller than the risk of untreated water - e.g. a cholera epidemic which could kill thousands

24
Q

What happens to the reducing power of halides going down group 7?

A

To reduce something, the halide ion needs to lose an electron from it’s outershell. Going down the group this attraction decreases because:
1. The ions get bigger, so the electrons are further away from the positive nucleus
2. There are extra inner electron shells, so there’s a greater shielding effect

Therefore the reducing power of halides increases down the group

25
Q

How does the reducing power of the halides effect the product of the next reaction formed after they have reacted with sulfuric acid?

A

All halides react with sulfuric acid to produce a hydrogen halide, after this they produce different products

26
Q

Give and explain the reaction between sodium fluoride and sulfuric acid

A

NaF(s) + H2SO4(l) -> NaHSO4(s) + HF(g)

  1. Hydrogen fluoride gas is formed. You’ll see misty fumes as the gas come into contact with the moisture on the air
  2. HF isn’t a strong enough reducing agent to reduce the sulfuric acid, so the reaction stops there
  3. It’s not a redox reaction - the oxidation states of the halide and sulfer stay the same (-1 and +6)
27
Q

Give and explain the reaction between sodium chloride and sulfuric acid

A

NaCl(s) + H2SO4(l) -> NaHSO4(s) + HCl(g)
​
1. Hydrogen chloride gas is formed. You’ll see misty fumes as the gas come into contact with the moisture on the air
2. HCl isn’t a strong enough reducing agent to reduce the sulfuric acid, so the reaction stops there
3. It’s not a redox reaction - the oxidation states of the halide and sulfer stay the same (-1 and +6)

28
Q

Give and explain the reaction between sodium bromide and sulfuric acid

A

NaBr(s) + H2SO4(l) -> NaHSO4(s) + HBr(g)

  • The first reaction gives misty fumes of hydrogenbromide gas
  • But the HBr is a stronger reducing agent than HCl and reacts with the H2SO4 in a redox reaction

2HBr(g) + H2SO4(l) -> Br2(g) + SO2 + 2H2O(l)

The reaction produces choking fumes of SO2 and orange fumes of Br2

Ox. state of S goes from +6 to +4 (reduction)

Ox. state of Br goes from -1 to 0 (oxidation)

29
Q

Give and explain the reaction between sodium iodide and sulfuric acid

A

NaI(s) + H2SO4(l) -> NaHSO4(s) + HI(g)

  • Same initial reaction giving HI gas
  • HI then reduces H2SO4 (like HBr):

2HI(g) + H2SO4(l) -> I2(g) + SO2 + 2H2O(l)

Ox. state of S goes from +6 to +4 (reduction)
Ox. state of I goes from -1 to 0 (oxidation)

  • But HI (being a very strong reducing agent) keeps going and reduces the SO2 to H2S

6HI(g) + SO2(g) -> H2S(g) + 3I2(s) + 2H2O(l)

Ox. state of S goes from +4 to -2 (reduction)
Ox. state of I goes from -1 to 0 (oxidation)

  • Solid iodine is also produced by this reaction

H2S gas is toxic

30
Q

What is the test for halide ions?

A
  1. Add dilute nitric acid to the solution to remove ions that may interfere with the test
  2. Add a few drops of silver nitrate solution
  3. A preciptate of silver nitrate will form

Ag(aq)+ + X(aq)- -> AgX(s)

31
Q

What happens when you test a solution containing fluoride ions using nitric acid?

A

No precipitate

32
Q

What happens when you test a solution containing chloride ions using nitric acid?

A

White precipitate forms

ppt. forms slowest of all the halogens

33
Q

What happens when you test a solution containing bromide ions using nitric acid?

A

Cream precipitate forms

34
Q

What happens when you test a solution containing iodide ions using nitric acid?

A

Yellow precipitate forms

35
Q

How can you confirm the results of the nitric acid test for halide ions

A

Add ammonia solution

36
Q

What happens to the precipitate formed from a nitric acid test to a solution that contained chloride ions when ammonia is added?

A

The white precipitate dissolves in dilute NH3(aq)

37
Q

What happens to the precipitate formed from a nitric acid test to a solution that contained bromide ions when ammonia is added?

A

The cream precipitate dissolves in concentrated NH3(aq)

38
Q

What happens to the precipitate formed from a nitric acid test to a solution that contained iodide ions when ammonia is added?

A

The yellow precipitate is insoluble in concentrated NH3(aq), therefore doesn’t redissolve

39
Q

Explain the flame test for Group 2 cations

A
  1. Dip a nichrome wire loop in conc. hydrochloric acid
  2. Then dip the wire into the unknown compound
  3. Hold the loop in the clear blue part of the bunsen burner flame
  4. Observe the colour change in the flame
40
Q

What colour can be observed when the Ca2+ (calcium) ion is tested using the flame test?

A

Brick red flame

41
Q

What colour can be observed when the Sr2+ (strontium) ion is tested using the flame test?

42
Q

What colour can be observed when the Ba2+ (barium) ion is tested using the flame test?

A

Pale green flame

43
Q

Explain the sodium hydroxide test for Group 2 cations

A
  • Add dilute NaOH dropwise to a test tube containing the metal ion solution and observe the precipitate that forms
  • Keep adding NaOH until it’s in excess
44
Q

What happens to a solution containing Mg2+ (magnesium) ions when dilute NaOH is added dropwise?

A

Slight white precipitate

45
Q

What happens to a solution containing Mg2+ (magnesium) ions when dilute NaOH is added in excess?

A

White precipitate

46
Q

What happens to a solution containing Ca2+ (calcium) ions when dilute NaOH is added dropwise?

A

Slight white precipitate

47
Q

What happens to a solution containing Ca2+ (calcium) ions when dilute NaOH is added in excess?

A

Slight white precipitate

48
Q

What happens to a solution containing Sr2+ (strontium) ions when dilute NaOH is added dropwise?

A

Slight white precipitate

49
Q

What happens to a solution containing Sr2+ (strontium) ions when dilute NaOH is added in excess?

A

Slight white precipitate

50
Q

What happens to a solution containing Ba2+ (barium) ions when dilute NaOH is added dropwise?

51
Q

What happens to a solution containing Ba2+ (barium) ions when dilute NaOH is added in excess?

52
Q

Explain the red litmus paper and NaOH test for ammonium ions

A
  • Ammonia gas is alkaline - so when exposed to a damp (damp so ammonia can dissolve) piece of litmus paper, the paper will turn blue
  • If hydroxide ion are added to a solution containing ammonium ions they will react to produce ammonia gas:

NH4+(aq) + OH-(aq)) -> NH3(g) + H2O(l)

Therefore you can use this to test whether a substance contains ammonium ions, add some dilute NaOH solution to the unknown substance in a test tube and gently heat the mixture, ammonia gas will be given off if positive for ammonium ions, confirm this using red litmus

53
Q

Explain how you test for sulfates with HCl and barium chloride

A
  • To identify a sulfate ion (SO42-), add a small amount of dilute hydrochloric acid, to remove and trace of carbonate ions
  • Followed by barium chloride solution (BaCl2(aq)

Ba2+(aq) + SO42-(aq) -> BaSO4(s)

  • If a white precipitate of barium sulfate is formed, it mean the origional compound contained a sulfate
54
Q

Explain how to test for hydroxide ions using litmus paper

A

Hydroxide ions make solutions alkali. Therefore to test if a solution contains OH- ions you can use a pH indicator e.g. red litmus paper

  1. Dip a piece of red litmus paper into the solution
  2. If hydroxide ions are present the paper will go blue
55
Q

Explain how hydrochloric acid can help to test for carbonates

A
  • When dilute hydrochloric acid is added to a solution conatianing carbonate ions, the solution will fizz
  • This is because the cabonate ions react with the hydrogen ions in the acid to produce carbon dioxide

CO32-(s) + 2H+(aq) -> CO2(g) + H2O(l)

  • The gas can then be tested using limewater for CO2
  • Carbon dioxide turn the limewater cloudy, so when bubbled into a test tube containing lime water, if the limewater goes cloudy the solution contained carbonate ions

3.2 A level Inorganic Chemistry (6 decks)

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