topic 3: group 7, the halogens Flashcards
trends in electronegativity in group 7
DECREASES DOWN GROUP
-As the atomic radius increases, the power of an atom to attract a pair of electrons decreases.
-This is because the atom is larger so there is less attraction from the electrons in the covalent bond to the nucleus.
-It is also being shielded more so this reduces the attraction even more.
trends in boiling points in group 7
INCREASE DOWN GROUP
-They are SIMPLE COVALENT MOLECULES
They are held together with VAN DER WAALS FORCES
-The strength of the van der waals forces increases as atomic radius increases and as molecule gets bigger as there are more electrons
-More van der waals forces take more energy to overcome
trends in oxidising ability in group 7
The oxidising ability of an atom is its ability to accept electrons from another element.
DECREASES DOWN THE GROUP
-Greater atomic radius and shielding.
-Less attraction between the outermost energy level and the positive nucleus, So harder to attract electrons.
-This means a more reactive (smaller) halogen will displace a less reactive halogen from its compound
-They are good oxidising agents as they accept electrons from the species being oxidised.
They themselves are being reduced.
Cl2 +2e- 🡪 2Cl-
displacement reactions of halide ions in aqueous solution
displacement reaction: A smaller halogen (more reactive, better oxidising agent) will displace a larger halide from its compound when a metal halide and halogen react.
-The REMOVED HALIDE ION WILL BE OXIDISED as its oxidation state will change from -1 to 0
-The MOST POWERFUL OXIDISING AGENT WILL DISPLACE THE LESS POWERFUL OXIDISING AGENT
-The metal halide solutions are colourless, Therefore there will be an observable change when any halogen is reacted with a metal halide solution. The observed COLOUR will show which halogen has been DISPLACED and in elemental form.
Cl2 + NaBr → Br2 + NaCl
Cl2 + 2Br - → 2Cl- + Br2
0 -1 -1 0
what is the reducing ability of an ion?
-The reducing ability of an ion is the ability of that ion to donate electrons to another element, reducing that element and causing that ion to itself be oxidised (losing electrons).
trend in reducing ability of the halide ions in group 7
REDUCING POWER INCREASES DOWN THE GROUP - BETTER REDUCING AGENTS
-Electrons are easier to lose from a larger ion with increased shielding
-weaker electrostatic attraction between the positive nucleus and the outermost electron as this electron becomes further from the positive nucleus
-easier for electrons to be donated from the halide ions
-Greater reducing power = greater reactivity and faster reduction
-Halide ions have a -1 charge
-In reactions they lose electrons and become halogen molecules
-In these reactions they are being oxidised.
I- > Br- > Cl-
Iodine ions are the best reducing agents as they lose electrons the easiest.
reaction of sodium halides with concentrated sulfuric acid
-the trend in reducing power can be seen when you react sodium halides with concentrated sulfuric acid
-the halide ions are in the sodium halides
reaction of sodium chloride with concentrated sulfuric acid
+1 -1 +1 +6 -2 +1 +1 +6 -2 +1 -1
NaCl(s) + H2SO4 (aq) → NaHSO4 (aq) + HCl (g)
This is an acid-base reaction, not a redox reaction.
All oxidation numbers are staying the same.
Only protons (hydrogens) are exchanged.
Chloride ions are not strong enough to reduce sulfur so the reaction stops.
OBSERVATIONS
HCl - steamy fumes
reaction of sodium bromide with concentrated sulfuric acid
+1 -1 +1 +6 -2 +1 +1 +6 -2 +1 -1
NaBr(s) + H2SO4 (aq) → NaHSO4 (aq) + HBr (g)
This is an ACID BASE REACTION, not a redox reaction.
The reaction doesn’t finish here though as Bromide ions are STRONG enough to REDUCE sulfuric acid into sulfur dioxide in a redox reaction.
+1 - 1 +1 +6 -2 0 +4 -2 +1 -2
2H+ + 2Br- + H2SO4 (aq) → Br2 (g) + SO2 (g) + 2H2O (l)
This is a REDOX reaction
Bromine has been OXIDISED from -1 to 0
It has REDUCED sulfur from +6 to +4
OBSERVATIONS
HBr - steamy fumes
Br2 - brown fumes and orange colour the solution when is dissolves
reaction of sodium iodide with concentrated sulfuric acid
1 -1 +1 +6 -2 +1 +1 +6 -2 +1 -1
NaI(s) + H2SO4 (aq) → NaHSO4 (aq) + HI(g)
This is an ACID BASE REACTION, not a redox reaction.
The reaction doesn’t finish here though as iodide ions are strong enough to reduce sulfuric acid into sulfur dioxide in a redox reaction.
+1 - 1 +1 +6 -2 0 +4 -2 +1 -2
8H+ + 8I+(g) + H2SO4 (aq) → 4I2 (s) + H2S (g) + 4H2O (l)
This is a REDOX reaction
Iodine has been oxidised from -1 to 0
It has reduced sulfur from +6 to -2
This is a very GOOD REDUCING AGENT, caused a large reduction.
OBSERVATIONS
HI - steamy fumes
I - solid black iodine
H2S - rotten egg smell
S - some yellow sulfur solid because it has been reduced past 0 (elemental oxidation state).
identifying aqueous metal halide ions using acidified silver nitrate
-All metal halides (except fluorides) react with silver ions in aqueous solution:
Cl-(aq) + Ag+(aq) 🡪 AgCl(s)
-Silver fluoride DOES NOT FORM A PRECIPITATE because it is SOLUBLE in water.
1) DILUTE NITRIC ACID HNO3(aq) is first added to the halide solution to REMOVE any SOLUBLE CARBONATE OR HYDROXIDE IONS as these would INTERFERE with the test by forming insoluble silver carbonate or hydroxide.
2) A few drops of SILVER NITRATE SOLUTION AgNO3(aq) are added and the halide precipitate forms
3) The colours of silver bromide and silver iodide are similar but we can distinguish between them by adding a few drops of CONCENTRATED AMMONIA SOLUTION NH3(aq). Silver bromide dissolves, silver iodide does not.
4) Solubility of the silver halide precipitates decreases down the group.
reaction of chlorine with water to form chloride ions and chlorate ions
Chlorine reacts with water in a reversible reaction to form chloric(l) acid, HCIO, and hydrochloric acid, HCl:
0 +1 -1
Cl2 (g) + H2O(l) ⇌ HClO(aq) + HCl(aq)
DISPROPORTIONATION reaction: chlorine is both oxidised and reduced.
-This reaction takes place when chlorine is used to purify water for drinking and in swimming baths, to prevent life-threatening diseases. Chloric(1) acid is an oxidising agent and kills bacteria by oxidation. It is also a bleach.
The other halogens react similarly, but much more slowly going down the group.
reaction of chlorine with water in sunlight to form chloride ions
2Cl2 (g) + 2H2O(l) → 4HCl(aq) + O2(aq)
Green colourless
This reaction is direct chlorination and is not reversible meaning in swimming pools you need to keep adding chlorine in to kill bacteria.
You can use an alternative reaction instead of direct chlorination which uses sodium chlorate which dissolves in water to form chloric acid. It is reversible.
NaClO(s) + H2O(l) ⇌ Na+(aq) + OH-(aq) + HClO(aq)
This reaction produces an alkaline solution, which shifts equilibrium to the left. To prevent this happening, swimming pools need to be kept slightly acidic so equilibrium moves right to make chloric acid. The OH- ion and the H+ (from acid) react to produce water, so this removes the product so equilibrium shifts to that side to counteract it.
However, this is carefully monitored and the water never gets acidic enough to corrode metal components and affect swimmers.
