Aqa A Level Chemistry - Halogens (Group 7)

Front 1

State the colour and physical state of fluorine

Back 1

Pale yellow

Gas

Front 2

State the colour and physical state of chlorine

Back 2

Green

Gas

Front 3

State the colour and physical state of bromine

Back 3

Red/brown

Liquid

Front 4

State the colour and physical state of iodine

Back 4

Grey

Solid

Front 5

State and explain what happens to the boiling points of the halogens as you go down group 7

Back 5

Boiling point increases

due to increased number of electrons

so strong van der Waals forces form between molecules

which require lots of energy to overcome

Front 6

State and explain what happens to the electronegativity of the halogens as you go down group 7

Back 6

Electronegativity decreases

due to a large distance between outer electron and nucleus

and shielding from electron shells

causing electrons to be less attracted to the nucleus

so halogen will be less able to attract electron pair in covalent bond

Front 7

State and explain what happens to the reactivity/oxidising power of the halogens as you go down group 7

Back 7

Reactivity decreases

due to outer shell being further away from the nucleus

so electrons are less attracted to the nucleus

Front 8

When will a halogen displace a halide from solution?

Back 8

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

Front 9

State the half equation when bromine water is added to potassium iodide

Back 9

Br2 + 2I ¯→ 2Br¯ + I2

Front 10

State and explain what happens when you add bromine water to potassium chloride

Back 10

Nothing happens

since chlorine is more reactive than bromine

so bromine can't displace it

Front 11

State and explain what happens when you add chlorine water to potassium bromide

Back 11

Orange solution formed

due to chlorine displacing bromide ions

since chlorine is more reactive that bromine

Front 12

State the equation for the formation of bleach

Back 12

2NaOH + Cl2 → NaClO + NaCl + H2O

Front 13

What is disproportionation?

Back 13

When a molecule is both oxidised and reduced in a reaction

Front 14

What type of reaction happens when chlorine reacts with water?

Back 14

A disproportionation reaction

Front 15

State the equation for the reaction of chlorine with water (not in sunlight)

Back 15

Cl2+ H2O ↔ HClO + HCl

Front 16

State the equation for the reaction of chlorine in sunlight

Back 16

2Cl2 + 2H2O ↔ 4HCl + O2

Front 17

State the advantages of treating water with chlorine

Back 17

Kills bacteria

Persists in the water to kill bacteria further down the supply

Stops algae growth, preventing bad tastes, smells and discolouration

Front 18

State the disadvantages of treating water with chlorine

Back 18

Chlorine gas is toxic

Chlorine as a liquid causes chemical burns

Chlorine may react with organic compounds from plants, forming carcinogenic chlorinated hydrocarbons

Front 19

State and explain what happens to the reducing power of halides as you go down the group

Back 19

Reducing power increases

due to less attraction between the outer electrons and the nucleus

caused by greater shielding and distance between the nucleus and outer shell

so electrons can easily be lost

Front 20

State the equation for the reaction between sodium fluoride and sulphuric acid, and state what you would observe

Back 20

NaF + H2SO4 → NaHSO4 + HF

Misty white fumes produced

Front 21

State the equations for the reactions between sodium bromide and sulphuric acid, and state what you would observe

Back 21

NaBr + H2SO4 → NaHSO4 + HBr

then

2HBr + H2SO4 → Br2 + SO2 + 2H2O

Misty white fumes of HBr

Choking fumes of SO2

Orange fumes of Br2

Front 22

State the equations for the reactions between sodium iodide and sulphuric acid, and state what you would observe

Back 22

NaI + H2SO4 → NaHSO4 +HI

then

2HI + H2SO4 → I2 + SO2 + 2H2O

then

6HI + SO2 → H2S + 3I2 + 2H2O

Solid, black I2 formed

Toxic H2S smells like bad eggs

Front 23

What safety precautions would you use when reacting a halide with sulphuric acid?

Back 23

Use a fume cupboard

Front 24

State the half equation for the formation of bromine ions and name the type of reaction taking place

Back 24

Br2 + 2e¯ → 2Br¯

Reduction

Front 25

Describe and explain how you would test for halides

Back 25

Add dilute nitric acid

to remove unwanted ions

then add a few drops of silver nitrate

to form a precipitate

Front 26

Give the general half equation for the reaction of silver nitrate with a halide

Back 26

Ag⁺ + X¯ → AgX

(where X is a halogen)

Front 27

How would you test a silver halide to determine which halide is present?

Back 27

Add ammonia solution and observing how they dissolve

Front 28

State the colours of the silver halide precipitates

Back 28

Fluoride: no precipitate

Chloride: white precipitate

Bromide: cream precipitate

Iodide: yellow precipitate

Front 29

State the solubility of the silver halides in ammonia solutions

Back 29

Chloride: white precipitate dissolves in dilute ammonia

Bromide: cream precipitate dissolves in concentrated ammonia

Iodide: yellow precipitate does not dissolve in ammonia

Front 30

Explain why iodide ions form a silver halide precipitate faster than the other halides.

Back 30

Iodide is a strong reducing agent

so its outer electrons are less attracted to its nucleus

since there's greater distance between its outer electrons and nucleus, and extra shielding

so it can easily lose its electrons

causing it to quickly reduce the silver nitrate