Group 7, Halogens

fluorine colour and state at room temperature

pale yellow gas

chlorine colour and state at room temperature

pale green gas

bromine colour and state at room temperature

brown-orange liquid

iodine colour and state at room temperature

grey solid

trend in boiling point down group 7

increases down the group

why does boiling point increase down group 7

van der waals forces increase due to increasing size and relative mass of atoms,

physical state goes from gas at the top of group 7 to solid at the bottom

trend in electronegativity down group 7

decreases down the group

define electronegativity

the power an atom has to attract a pair of electrons to itself from a covalent bond

why does electronegativity decrease down group 7

• atomic radius increases, distance between positive nucleus and bonding electrons increases = less attraction

• more shielding down group 7

what do more reactive halogens displace in displacement reactions

more reactive halogens will displace less reactive halide ions

trend in reactivity down group 7

decreases down group 7

why does reactivity decrease down group 7

• atomic radius increases, less ability to attract electrons compared to atoms with a smaller radius

trend in oxidising power down group 7

less oxidising as we go down,

shown by reacting halogens with halide ions

why does oxidising power decrease down group 7

atoms with a smaller radius attract electrons easier,

for a reaction to occur, an electron must be gained,

atomic radius increases down group 7, so harder to attract an electron

when will a halogen displace a halide from solution

if the halide is LOWER in the periodic table,

more reactive halogens displace less reactive halides

in terms of electrons, what must occur for a displacement reaction to occur

an electron must be gained

does displacement occur when chlorine water is added to potassium chloride solution

no reaction,

both have Cl which is the same reactivity

is there a displacement reaction when chlorine water is added to potassium bromide solution

yes,

Cl2 + 2Br → 2Cl- + Br2

what can be observed from the displacement reaction between addition of chlorine water (almost colourless) to potassium bromide solution (colourless)

orange solution, as Br2 is made

is there a displacement reaction between potassium iodide solution (KI, colourless) and chlorine water

yes

Cl2 + 2I- → 2Cl- + I2

what is observed during displacement reaction between potassium iodide solution and chlorine water

brown solution, as iodine (I2) is made

is there a displacement reaction between potassium chloride solicitor and addition of bromine water (orange)

no reaction, bromine is less reactive than chlorine

is there a displacement reaction between potassium bromide solution and bromine water

no reaction, both contain bromine which are the same reactivity

is there a displacement reaction between potassium iodide solution and bromine water? what is observed?

yes,

brown solution is formed as I2 is made

Br2 + 2I- → 2Br- + I2

bromine is more reactive than iodide

is there a displacement reaction between potassium chloride solution and addition of iodine solution (brown)

no reaction, iodine is less reactive/oxidising than chloride ion

is there a displacement reaction between potassium bromide solution and iodine solution

no reaction, iodine is less reactive/oxidising than bromide

is there a displacement reaction between potassium iodide solution and iodine solution

no reaction, both contain iodine which have the same reactivity

which reaction is bleach made from

disproportionation reaction

mixing chlorine and sodium hydroxide will form sodium chlorate (I) solution = bleach

what is a disproportionation reaction

a redox reaction where an element is both oxidised and reduced

reaction to make bleach?

chlorine + sodium hydroxide → sodium chlorate (I)

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

which element has been reduced and oxidised in the reaction to make bleach and how

chlorine

reactant: Cl2, oxidation state of 0

product: NaClO, chlorine has an oxidation state of +1

product: NaCl, chlorine has an oxidation state of -1

uses of sodium chlorate (I), bleach

treating water

bleach paper and fabrics

cleaning reagent (bleach)

importance of water sterilisation

stop outbreaks of water borne diseases

what do we add to water to sterilise it

chlorine

how does adding chlorine to water sterilise it

chlorate (I) ions (ClO-) kill bacteria,

useful in drinking water and pools

reaction between water and chlorine

H2O (l) + Cl2 (g) → 2H+ (aq) + Cl- (aq) + ClO- (aq)

how is the reaction between chlorine and water a disproportionation reaction

chlorine is oxidised and reduced

what can decompose chlorinated water

sunlight

reaction between water and chlorine in strong sunlight

2H2O (l) + Cl2 (g) → 4H+ (aq) + 2Cl- (aq) + O2 (g)

issue with sunlight decomposing water and chlorine? how is this solved?

no ClO- is made, so bacteria is not killed (no active ingredient)

in swimming pools, disease can spread as sunlight is concentrated

so we constantly top up the chlorine in swimming pools

advantages of chlorinating drinking water

• destroys microorganisms which cause disease

• long lasting so reduce bacteria build up further down the supply

• reduces growth of algae that discolours water and can give it a bad smell and taste

• small amounts of chlorine is used, so risk of cancer is very LOW

disadvantages of chlorinating drinking water

• chlorine gas is toxic and irritates the respiratory system, in high concentrations

• liquid chlorine causes severe chemical burns to the skin, in high concentrations

• chlorine can react with organic compounds present in water to make chloroalkanes - have links to causing cancer, but risk of not chlorinating water could lead to a cholera epidemic

how are halide ions reducing agents

they lose an electron in reactions

define reducing agent

oxidised themselves, lose electrons

trend in reducing power of halide ions down the group

reducing power increases down the group

F- , weakest

Cl-

Br-

I-, strongest

why does reducing power of halides increase down group 7

• ionic radius increases

• distance between nucleus and outer electrons become larger and there’s more shielding, weakening the attractive force

• outer electron is lost more readily (reducing agents lose an electron during reactions)

2 tests to prove reducing power of halide ions

1 reaction with sulfuric acid

2 reaction with silver nitrate

reaction between sodium chloride and sulfuric acid

H2SO4 + NaCl → NaHSO4 + HCl

what is observed from the reaction between sulfuric acid and sodium chloride

white, steamy fumes, from HCl gas produced

why isn’t the reaction between sodium chloride and sulfuric acid a redox reaction

oxidation state of sulfur remains the same, +6

so, NaHSO4 is not a reduction product

why is there no further reaction between chloride ions and sulfuric acid after production of NaHSO4

Cl- is not a strong enough reducing agent to reduce sulfur

reduction products of sulfur

SO2 (+4)

S (0)

H2S (-2)

first reaction of sodium bromide with sulfuric acid

HS2SO4 + NaBr → NaHSO4 + HBr (displacement, not redox)

second/redox reaction between bromide ions and sulfuric acid, using products from the first x

2HBr + H2SO4 → Br2 + SO2 + 2H2O

what can be observed during the reduction of sulfur by bromide ions

Br2 gas produced = orange vapour

what is oxidised and reduced in the reaction between HBr and H2SO4, state oxidation state

Br- ions are oxidised

Sulfur in H2SO4 (+6) is reduced to SO2 (+4)

first reaction between sodium iodide with sulfuric acid

NaI + H2SO4 → NaHSO4 + HI (displacement)

second reaction of hydrogen iodide, HI, with sulfuric acid

2HI + H2SO4 → I2 + SO2 + 2H2O

what is oxidised and reduced between the reaction of hydrogen iodide and sulfuric acid

iodide ions are oxidised to I2

sulfur in H2SO4 (+6) is reduced to SO2 (+4)

further reduction equation between hydrogen iodide and sulfuric acid

6HI + H2SO4 → 3I2 + S + 4H2O

H2SO4 (+6) reduced to S (0)

what can be observed during this further reduction reaction of sulfuric acid

yellow solid, as sulfur (s) is produced

full reduction of sulfuric acid with hydrogen iodide/iodide ions

other equations here, then:

8HI + H2SO4 → 4HI2 + H2S + 4H2O

what is observed/smelt from the full reduction of sulfur using iodide ions

rotten egg smell, from H2S (g)

why does greater reducing power = longer reactions

halide is powerful enough to reduce more species

how to test for halides

1 add dilute nitric acid (HNO3)

2 add silver nitrate

3 confirm halide with ammonia solution

why do we add dilute nitric acid when testing for halides with silver nitrate

to remove any carbonite/sulfite impurities, so there’s no false result

simple ionic equation between Ag+ and halides

Ag+ (aq) + Cl- (aq) → AgCl (s)

Ag+ (aq) + Br- (aq) → AgBr (s)

Ag+ (aq) + I- (aq) → AgI (s)

what do these reactions produce, why is a precipitate formed

produce INSOLUBLE silver halides

halide test results after adding silver nitrate

AgF / F- = none

AgCl / Cl- = white

AgBr / Br- = cream

AgI / I- = yellow

why doesn’t AgF produce a precipitate

AgF is soluble

how do further tests confirm the halide present

(results when ammonia is added)

add ammonia,

• AgCl dissolves in dilute ammonia

• AgBr dissolves in concentrated ammonia

• AgI is insoluble in any concentration of ammonia

test for group 2 cations

flame test used for a solid sample,

1 dip nichrome wire into HCl to clean

2 dip into sample

3 place loop into blue flame and observe colour

results for positive group 2 cation test, Ca2+, Sr2+, Ba2+

Ca2+ = dark red

Sr2+ = red

Ba2+ = green

can flame test be used for solutions? issue with this?

yes, it can be made into a solution and dipped into the loop, but will be difficult if the sample is insoluble

how to test for ammonium compounds with litmus paper + positive result

1 add sodium hydroxide and gently heat

2 if ammonium compound is present, ammonia gas is produced (alkaline)

3 damp red litmus paper turns blue is alkaline ammonia is produced

reaction between ammonium ions and hydroxide ions?

test for hydroxides with litmus paper, issues with test?

hydroxides are alkaline so will turn red litmus paper blue,

this does not fully confirm a hydroxide because red litmus turns blue for any alkaline

test for carbonates

• HCl + carbonate → CO2 (g)

• bubble through limewater = cloudy if CO2 gas present

test for sulfates

• add HCl to remove any carbonates

• add barium chloride, BaCl2

• positive test = white precipitate

what forms the white precipitate in sulfate test

BaSO4 (s), barium sulfate - insoluble

Ba2+ (aq) + SO4 2- (Aq) → BaSO4 (S)

order of tests to prevent false positives

1 carbonates

2 sulfates

3 halides