Halogens Flashcards
Reasons for the trends in m.pt and b.pt for group 7 elements
Group 7 elements are simple covalent molecules, held together by London forces
moving down group 7, the number of electrons per atom increase, meaning the London foces increase in strength
Requiring more energy to overcome
Reasons for the trends in physical state at room temp of group 7 elements
- Fluorine and chlorine are gases at room temperature because they have weak London forces between molecules
- Bromine is a liquid because it has more electrons and slightly stronger London foces
- Iodine and astatine are solids at room temp because they have the most number of electrons in the group and therefore have stronger London foces between molecules
Reasons for the trends in electronegativity for group 7 elements
Moving down group 7, despite proton number increasing:
* Atomic radius increases, distance between nucleus and outer shell increases
* Number of shells (so shielding) increases
therefore the ability of the halogens to attract electrons decrease
Reasons for trends in reactivity of group 7 elements down the group
Moving down group 7:
* Atomic radius increases, distance between nucleus and outer shell increases
* Number of shells (so shielding) increases
therefore the ability of a halogen to attract an electron decreases
Oxidising power of halogens
Halogens are good oxidising agents
* they accept electrons from the species being oxidised and are reduced themselves
* oxidising power decreases down the group as ability to attract electrons decrease
* Relative oxidising strength means a halogen will displace any halide beneath it in the Periodic Table
Reducing power of halide ions
Halide ions are good reducing agents
* They donate electrons to the species being reduced and are oxidised themselves
* Reducing power increases down the group as electrons are more easily lost from larger ions
Trend in reactivity of group 7 elements - redox reactions of Cl2, Br2 and I2 with halide ions in aqueous solution, followed by the addition of an organic solvent
Oxidising power decreases moving down the group as ability to attract electrons decreases:
* Chlorine can displace bromide and iodide
* Bromine can displace iodide
* Iodide cannot displace either
Ex.
Cl₂ + 2Br- –> Br₂ + 2Cl-
Halogens oxidation reactions with group 1 and 2 metals (In terms of oxidation numbers)
Group 1 and 2 metals react with chlorine gas to form metal chlorides, which are white precipitates
Chlorine is reduced and metal is oxidised
Chlorine disproportionation reaction with water (In terms of oxidation numbers)
Chlorine reacts with cold water:
- produces chlorate (I) ions and (ClO-) and chloride ions
- chlorine is both oxidised and reduced, therefore disproportionation reaction
- oxidisation state goes from 0 to both -1 and +1
Cl2 + H2O –> ClO- + Cl- + 2H+
The use of chlorine in water treatment
The disproportionation reaction of water and chlorine is used to kill bacteria
This can pose risks as chlorine is toxic
Chlorine disproportionation reaction with cold, dilute aqueous sodium hydroxide to form bleach (In terms of oxidation numbers)
Chlorine reacts with cold dilute sodium hydroxide is used in bleach production
Sodium chlorate (a key ingredient in the production of bleach) is produced via the reaction
2NaOH + Cl2 –> NaClO + NaCl + H2O
Chlorine disproportionation reaction with hot alkali (In terms of oxidation numbers)
Chlorine reacts with hot, concentrated alkali to be disproportionated even further than thechlorine with cold dilute sodium hydroxide reaction
forms a species with oxidisation number -1 and +5 (chlorine)
3Cl2 + 6NaOH –> NaClO3 + 5NaCl + 3H2O
The reaction of solid Group 1 halides with concentrated sulfuric acid (to illustrate the trend in reducing ability of the hydrogen halides)
Redox reactions between Halides and H2SO4 vary depending on the reducing ability of the halide
The greater the reducing power, the further the reaction will proceed as the halide is powerful enough to reuce more species
Therefore reducing power increases down group 7
1. Fluoride and Chloride ions
- Both aren’t great reducing agents
- Only one reaction occurs, HF and HCl are misty fumes
NaF + H2SO4 –> NaHSO4 + HF
NaCl + H2SO4 –> NaHSO4 + HCl
2. Bromide ions
- HBr is observed as a misty fume
- Orange fumes of Br2 and choking fumes of SO2 are observed from the second reaction (where HBr reduces H2SO4)
NaBr + H2SO4 –> NaHSO4 +HBr
2HBr + H2SO4 –> Br2 + SO2 + 2H2O
3. Iodide ions
- HI is observed as misty fumes from the 1st reaction
- Hi then reduces to solid iodine and choking fumes of SO2
- HI then further reduces SO2 to toxic gas H2S (which smells of bad eggs! stinky!)
NaI + H2SO4 –> NaHSO4 + HI
2HI + H2SO4 –> I2 + SO2 + 2H2O
6HI + SO2 –> H2S + 3I2 + 2H2O
Precipitation reactions of the aqueous anions Cl-. Br- and I- with aqueous silver nitrate solution
- Acidified silver nitrate solution can be used to distinguish between halide ions
(The silver nitrate solution must be acidified using dilute nitric acid first, to remove any excess ions present in the solution that might react and affect results)
- A precipitate of the silver halide will form, the colour of which helps identify the halide ion present
General equation, where X is a halogen:
Ag+(aq) + X-(Aq) –> AgX(s)
Halide ion, colour of precipitate formed with silver:
* Fluoride - No precipitate
* Chloride - White precipitate
* Bromide - Cream precipitate
* Iodide - Yellow precipitate
using aqueous ammonia solution to differentiate silver halide precipitates
Different silver halides have different solubilities in ammoniaD, therefore ammonia can be used to distinguish between halides.
Silver halide, Solubility of halide in ammonia:
* Silver Chloride - dissolves in dilute ammonia
* Silver Bromide - dissolves in concentrated ammonia
* Silver Iodide - Insoluble in concentrated ammonia
