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AS chemistry unit 1 > 1.8 halogen > Flashcards

1.8 halogen Flashcards

1
Q

fluorine

A

F2 yellow gas

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2
Q

chlorine

A

Cl2 yellow-green gas

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3
Q

bromine

A

Br2 red-brown Liquid*
(* red-brown vapour when heated)

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4
Q

iodine

A

I2 grey-black solid (purple vapor when heated)

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5
Q

astatine

A

At2 Black solid

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6
Q

what does iodine uniquely do

A

Note that liquid bromine vaporises readily at room temperature and iodine sublimes
(changes from solid to gas without passing through the liquid state) when heated to
give a purple vapour.

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7
Q

solubility

A

The halogens are non-polar molecules and do not dissolve well in polar solvents eg
water or ethanol. They dissolve much more readily in non-polar solvents such as
hexane because they both exhibit the same main intermolecular force, namely van
der Waals’ attractions. The solubility of the halogens can be explained by using the
principle that ‘like-dissolves-like’
. Therefore, iodine, which is non-polar and
molecular covalent, is more soluble in hexane which is also non-polar and molecular
covalent rather than in water, which is polar.

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8
Q

Determine the solubility of chlorine and iodine in aqueous and non-aqueous solvents

A

Method
• add 1 cm3 of each of the aqueous halogen solutions (chlorine water, iodine
solution) into separate test tubes.
• add equal volumes of a saturated hydrocarbon solvent e.g. hexane to each test
tube, stopper the test tube and shake the mixture by inverting the test tube a few
times.
• allow the two layers to settle; observe and record the colour

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9
Q

halogen solubility in water

A

Chlorine Soluble, forming a pale
green/colourless solution
(‘chlorine water’)

Bromine Soluble, forming a
orange/yellow/brown
solution

Iodine Virtually insoluble, but any
solution formed is
brown/yellow

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10
Q

halogen solubility in hexane

A

Chlorine Soluble, forming a pale
green/colourless solution

Bromine Soluble, forming a orange
solution

Iodine Soluble, forming a purple
solution

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11
Q

reaction of halogen with aqueous alkali

A

Conditions: cold (approx 10oC), dilute NaOH

Cl2 (g) + 2NaOH(aq)  NaClO(aq) + NaCl(aq) + H2O(l)
sodium hypochlorite/
sodium chlorate(I)
The general ionic equation for this reaction is:
Cl2 (g) + 2OH-
(aq) ->ClO-
(aq) + Cl-
(aq) + H2O(l)
NaClO (bleach) is called sodium hypochlorite or sodium chlorate(I) as it contains the
chlorate (I) ion, ClO-
. The oxidation number of chlorine in chlorate (I) is +1.

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12
Q

The chlorate(I) ion can decompose in two ways:

A
  1. On exposure to sunlight, oxygen is evolved:
  2. On warming the solution, the ion disproportionates:
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13
Q

Conditions: hot (at or above 70oC), concentrated NaOH

A

3NaClO (aq)  2NaCl(aq) + NaClO3 (aq)
The overall equation for the reaction of chlorine with HOT aqueous alkali, such as
NaOH, can be represented as:
3Cl2 + 6NaOH  NaClO3 + 5NaCl + 3H2O
The general ionic equation for this reaction is:
3Cl2 + 6OH-  ClO3
- + 5Cl- + 3H2O

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14
Q

chlorine with water

A

Chlorine reacts slowly with water over time producing a mixture of two acids:
hydrochloric acid (HCl) and hypochlorous acid (HClO), (systematic name chloric(I)
acid). There will be chloride ions (Cl-) and chlorate(I) ions (ClO-) in the solution
produced, as well as hydrogen ions which will make chlorine water slightly aci

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15
Q

Adding Chlorine / Ozone to Drinking Water

A

Chlorine and ozone may both be added to drinking water to kill bacteria and sterilise
it, making it safer to drink.
Ozone, O3, is an allotrope of molecular oxygen, O2. It can be generated by a high
voltage electric discharge through oxygen, or on a smaller scale using ultaviolet
light.
3O2 → 2O3
Chlorine reacts with water to form a mixture of hypochlorous acid (HClO) and
hydrochloric acid (HCl):
Cl2 + H2O → HClO + HCl

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16
Q

Both O3 and HClO kill bacteria by oxidising the microorganisms.
The table below details some of the advantages and disadvantages of using chlorine
and ozone in the treatment of water:

A

chlorine; Advantages
Chlorine Cheaper than ozone (cost
effective)
Provides residual protection as
it is still present in water when
it reaches the consumer
More soluble in water than
ozone
disadvantage
Cannot kill some
microorganisms
Leaves chemicals in the water
Unpleasant taste
Toxic to humans, so very low
quantities used
Ozone ; advantages
More effective at killing
bacteria and kills more
different types
Breakdown product is oxygen
No residual chemicals left in
the water
Reacts with natural organic
matter far better than chlorine
does and removes it from
water
No unpleasant smell or taste
disadvantages
More expensive than chlorine –
equipment and operational costs
Does not provide residual
protection against
microorganisms
Less soluble in water than
chlorine so requires special
mixing techniques
There are potential fire hazards
and toxicity issues associated
with ozone generation

17
Q

Explain why the public water supply may be fluoridated and why some people are
opposed to this

A

Fluoride is added to water to reduce tooth decay. The fluoridation of water is
opposed on the grounds that individuals do not have the right to choose.

18
Q

Oxidising ability of the halogens

A

All halogens are oxidising agents as they can accept electrons and are reduced.
F2 + 2e-  2F-
The oxidising power of the halogens decreases as atomic number increases from
fluorine to chlorine to bromine to iodine (as you move down the group).
Oxidising agents readily accept electrons; the ease with which it gains electrons
determining how effective it is as an oxidising agent.

19
Q

displacement reactions

A

From our knowledge of the oxidising powers of the halogen (see previous section)
we would expect fluorine to displace all other halides from a solution of a halide
compound. This is because the halide ions are oxidised by a better oxidising
agent (halogen higher in the group) to form the corresponding halogen. We
would also expect chlorine to displace bromide and iodide from solutions of halide
compounds, and bromine to displace iodide in solution.
A more reactive halogen will displace a less reactive halogen from a
compound.

20
Q

Produce a reactivity order of the halogens using the displacement reactions of
the halogens with other halide ions in solution

A

Method
• add 1 cm3 of each of the aqueous halogen solutions (chlorine water, bromine
water, iodine solution) into separate test tubes.
• add 1 cm3 of potassium chloride solution to each and record any colour change.
• repeat the experiment with potassium bromide solution and then with potassium
iodide solution.
• a saturated hydrocarbon e.g. hexane could be added to distinguish between
bromine and iodine or starch solution could be added to detect iodine.

21
Q

Reducing ability of the halides

A

The reducing power of the halide ions increases as atomic number increases from
fluoride to chloride to bromide to iodide (as you move down the group).
Reducing agents (also called reductants) cause a reduction to occur so they cause
another atom, molecule or ion to gain electrons. This means that a reducing agent
loses electrons and becomes oxidised.
Reducing agents readily donate electrons; the ease with which a reducing agent
loses electrons determines how effective it is as a reducing agent.
Halide ions lose electrons from their outer energy level.

22
Q

Reaction of the solid halides with conc. sulfuric acid

A

The trend in the reducing ability of the halide ions can be demonstrated in the
reaction between solid halide salts and concentrated sulfuric acid.
The hydrogen halides can be prepared by reacting the corresponding sodium halide,
NaX, with concentrated sulfuric acid, H2SO4, according to the following general
equation:
NaX + H2SO4 → NaHSO4 + HX Where x = F, Cl, Br or I
Note: sodium sulfate is not formed
Hydrogen fluoride and hydrogen chloride prepared by this method are pure.
However, as the oxidising ability of the halogens decreases down the group, the
corresponding halide (X-) becomes a better reducing agent to the extent that it
reduces the concentrated sulfuric acid and forms the corresponding halogen.

23
Q

Carry out the reactions of the halides with concentrated sulfuric and
phosphoric acids and perform chemical tests for the products (excluding
hydrogen sulfide)

A

Method
• place 1 spatula of solid potassium chloride in test tube.
• add a few drops of concentrated sulfuric acid and record observations.
• test the gas evolved.
• repeat the experiment using solid potassium bromide.
• repeat the experiment using solid potassium iodide.
• repeat the experiment using concentrated phosphoric acid instead of concentrated
sulfuric acid.

24
Q

fluoride with conc. sulfuric acid

A

Fluoride
NaF(s) + H2SO4  HF(g) + NaHSO4
misty/steamy fumes
This is an acid-base reaction; HF is not a sufficiently strong reducing agent (see
introduction above and previous section) to reduce sulfuric acid and so no redox
reaction takes place.
Observations: misty/steamy fumes (HF); heat released; gas produced; solid
disappears; pungent smell (HF)
Names of products: sodium hydrogensulfate and hydrogen fluoride

25
Q

chloride with conc sulfuric acid

A

Chloride
NaCl(s) + H2SO4  HCl(g) + NaHSO4
misty/steamy fumes
This is an acid-base reaction; HCl is not a sufficiently strong reducing agent to
reduce sulfuric acid and so no redox reaction takes place.
This reaction can be used as a method for preparing HCl. It may not be used to
prepare HBr or HI as these substances act as reducing agent and go on to reduce
the conc. sulfuric acid as soon as they are formed
Observations: misty/steamy fumes (HCl); heat released; gas produced; solid
disappears; pungent smell (HCl).
Names of products: sodium hydrogensulfate and hydrogen chloride

26
Q

bromide with conc sulfuric acid

A

Bromide
Sodium bromide reacts in a similar way to sodium chloride, producing misty fumes of
hydrogen bromide gas:
NaBr(s) + H2SO4  HBr(g) + NaHSO4
misty/steamy fumes
However, bromide ions are stronger reducing agents than chloride and fluoride ions
and after the initial acid-base reaction, they reduce the sulfur in H2SO4 from +6 to +4
in sulfur dioxide, SO2.
2HBr + H2SO4  Br2 + SO2 + 2H2O
red-brown vapour
This reaction must be carried out in a fume cupboard.
Observations – misty/steamy fumes (HBr); heat released; gas produced; solid
disappears; red-brown vapour (bromine); pungent smell (HBr/SO2/Br2).
Names of products: sodium hydrogensulfate; hydrogen bromide; bromine; sulfur
dioxide and water.

27
Q

iodide with conc. sulfuric acid

A

Iodide
When sodium iodide reacts with concentrated sulfuric acid, HI is formed. Iodide
ions are the strongest halide reducing agent. As a result, it can give eight
electrons to the sulfur – reducing it from +6 in sulfuric acid to -2 in hydrogen sulfide:
NaI + H2SO4  HI + NaHSO4
S = +6 steamy/misty
fumes of HI
2HI + H2SO4  I2 + SO2 + 2H2O
purple S = +4
vapour
6HI + H2SO4  3I2 + S + 4H2O
S = 0
8HI + H2SO4  4I2 + H2S + 4H2O
S = -2
Observations: misty/steamy fumes (HI), heat released, gas produced, solid
disappears, purple vapour (I2) and grey-black solid on sides of the test tube (iodine),
pungent smell (HI/SO2/I2) rotten egg smell (H2S), yellow solid (S)
Names of products: sodium hydrogen sulfate; hydrogen iodide; iodine; sulfur dioxide;
sulfur; hydrogen sulfide; water

28
Q

Solid halides with concentrated sulfuric acid OBSERVATIONS

A

Solid halides with concentrated sulfuric acid OBSERVATIONS
fluoride steamy/misty fumes of HF
chloride steamy/misty fumes of HCl
bromide
steamy/misty fumes of HBr
red-brown vapour Br2
iodide
steamy/misty fumes of HI
purple vapour (I2)
smell of rotten eggs (H2S)
yellow solid (S)
grey-black solid (I2)

29
Q

Solid halides with concentrated phosphoric acid OBSERVATIONS

A

Solid halides with concentrated phosphoric acid OBSERVATIONS
fluoride steamy/misty fumes of HF
chloride steamy/misty fumes of HCl
bromide steamy/misty fumes of HBr
iodide steamy/misty fumes of HI

30
Q

Halides and concentrated phosphoric acid

A

NaF(s) + H3PO4  HF(g) + NaH2PO4
NaCl(s) + H3PO4  HCl(g) + NaH2PO4
KBr(s) + H3PO4  HBr(g) + KH2PO4
KI(s) + H3PO4  HI(g) + KH2PO4

AS chemistry unit 1 (10 decks)

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