Classwork Series and Exercises {Chemistry – SS1}: Halogens

Chemistry SS 2 Week 1

Topic: Halogens

Introduction

The halogens are the collective name given to the elements in group VII of the Periodic Table. There are five halogens; fluorine, chlorine, bromine, iodine and astatine. Astatine is very radioactive and cannot exist for more than a few microseconds before decaying. We will thus be concerned with the chemistry of fluorine, chlorine, bromine and iodine.

All these elements are most commonly found in the -1 oxidation state, as X^– ions. These are known as halide ions.

Halide is the name given to the ion of halogens. Table below shows the corresponding halide of the halogen

Halogen	Halide
Fluorine	Fluoride
Chlorine	Chloride
Bromine	Bromide
Iodine	Iodide

Structure

Since each atom in this group has seven valence electrons, they tend to form diatomic molecules, e.g. F₂, Cl₂, Br₂ and I₂. They are thus simple molecular, with weak intermolecular forces between the molecules.

Appearance and Colour

Halogen	In pure form	In non-polar solvents	In water
Fluorine	Pale yellow gas	–

(Reacts with solvents)-

(Reacts with water) Chlorine Pale green gas Pale green solution

Pale green solution Bromine Dark red liquid Orange solution

Orange solution Iodine Grey solid Purple solution-

(Insoluble)

but forms a brown solution if excess KI is present

The halogens are usually used in aqueous solution. Although iodine is insoluble in water, it is soluble if iodide ions are present (the iodine reacts with iodide ions to form triiodide ions as follows: I₂(aq) + I^–(aq) à I₃^–(aq). The triiodide ions give the solution its brown colour.

Melting and Boiling Points

Halogen	Melting point /oC	Boiling point /oC
Fluorine	-220	-188
Chlorine	-101	-35
Bromine	-7	59
Iodine	114	184

The melting and boiling points of the halogens increase steadily down the group. This is due to the increase in strength of the Van Der Waal’s forces between the molecules, which results from the increasing number of electrons in the molecule and the increasing surface area of the molecule.

Electronegativity

Electronegativity is the ability of an atom to attract electrons in a covalent bond. The electronegativity of the halogen atoms decreases down a group:

	Halogen	Electronegativity
	F	4.0
	Cl	3.0
	Br	2.8
	I	2.5

As the number of shells increases, the shielding increases and the electrons in the covalent bond are further from (and more shielded from) the nucleus. Therefore they are less strongly attracted to the nucleus and the electronegativity decreases.

Electronic Configuration

Element	Electron arrangement of atom
Fluorine / F	2.7
Chlorine / Cl	2.8.7
Bromine / Br	2.8.18.7
Iodine / I	2.8.18.18.7
Astatine / At	2.8.18.32.18.7

Chemical tests

• Chlorine – Turns damp litmus red, and then bleaches it.
• Bromine – Bromine also turns damp litmus red and bleaches it but slower. When sodium hydroxide is added, bromine loses its colour.
• Iodine – When starch solution is added, a blue/black colour forms

Physical properties

• Non-metals
• Insulators of electricity
• Poor conductors of heat
• Low melting point and boiling point
• Melting point and boiling point increases (going down the group) – molecular size increases / forces of attraction between molecules become stronger & more heat energy is required to overcome the stronger forces of attraction (Van der Waals’ forces of attraction between molecules).
• Colour: F₂ (pale yellow gas), Cl₂ (greenish-yellow gas), Br₂ (reddish-brown liquid), I₂ (purplish-black solid) and At₂ (black solid – rarest naturally occurring element and extremely radioactive)
• Low density
• Density of element increases (going down the group) – increase in atomic mass

Chemical properties

• Oxidation – reduction behaviour

The halogens can act as oxidising agents by gaining electrons to form halide ions.

X₂ + 2e^– → 2X^–

The oxidising ability decreases down the group with fluorine being the strongest oxidising agent. In simple terms the formation of X^– occurs in two steps.

A strong oxidising agent will readily form X^–. It will therefore, have a low bond energy and a high electron affinity. Both bond energies and electron affinities decrease down the group as the nuclear pull on the outer electrons decrease.

The change in electron affinities is most significant and so oxidising ability decreases down the group as the nuclear pull on the outer electrons decreases down the group.

Fluorine has anomalous properties due to its small size. It has a lower electron affinity than Chlorine due to electron repulsions in the overcrowded outer shell. However, it is a stronger oxidising agent than Chlorine due to its unusually low bond energy. This is due to the repulsions between outer electrons in the F₂ molecule.

Halide ions act as reducing agents by losing electrons.

2X^– – 2e^– → X₂

The reducing ability increases from F^– to I^–.

In all their reactions the halogens act as oxidising agents. Hence, their reactivity decreases down the group.

• Reaction with water

F₂ oxidises H₂O to O₂ gas in a very exothermic reaction.

2F_2(g) + 2H₂O_(l) → O_2(g) + 4HF_(g)

Cl₂ dissolves in H₂O and some hydrolysis occurs. A yellow solution of ‘chlorine water’ is formed which is a mixture of two acids. No O₂ is evolved.

Cl_2(g) + H₂O_(l) → HCl_(aq) + HOCl_(aq)

Br is only slightly soluble in H₂O and there is less hydrolysis.

Br_2(l) + H₂O_(l) → HBr_(aq) + HOBr_(aq)

I₂ is virtually insoluble in H₂O. It is however soluble in KI solution due to the formation of the triiodide anion.

I_2(s) + I^–_(aq) → I₃^–_(aq)

Note: All halogens are more soluble in non-polar solvents such as CCl₄. Cl₂ gives a colourless solution. Br₂ a red solution and I₂ a violet one.

• Displacement reactions

If a more reactive halogen is placed into a solution containing a less reactive halide a displacement reaction is seen.

For example: Cl_2(g) + KI_(aq) → KCl_(aq) + I_2(s)

• F₂ will displace Cl^– Br^– I^–
• Cl₂ will displace Br^– I^–
• Br₂ will displace I^–
• I₂ will displace none

• Reaction with alkali ( NaOH solution)

This reaction depends on the conditions:

Cold NaOH solution.

2OH^–_(aq) + Cl_2(g) → Cl^–_(aq) + ClO^–_(aq) + H₂O_(l)

Hot NaOH solution.

6OH^–_(aq) + Cl_2(g) → 5Cl^–_(aq) + ClO₃^– + 3H₂O_(l)

In hot solution the ClO^– ions are disproportionating

3ClO^–_(aq) → 2Cl^–_(aq) + ClO₃^–_(aq)

• Other Reactions of the Halogens

Reaction with hydrogen H₂

• Halogens readily combine with hydrogen to form the hydrogen halides which are colourless gaseous covalent molecules
• e.g. hydrogen + chlorine ==> hydrogen chloride
• H_2(g) + Cl_2(g) ==> 2HCl_(g)
• The hydrogen halides dissolve in water to form very strong acids with solutions of pH1 e.g. hydrogen chloride forms hydrochloric acid in water HCl(aq) or H⁺Cl^–(aq) because they are fully ionised in aqueous solution even though the original hydrogen halides were covalent! An acid is a substance that forms H⁺ ions in water.
• Bromine forms hydrogen bromide gas HBr_(g), which dissolved in water forms hydrobromic acid HBr_(aq). Iodine forms hydrogen iodide gas HI_(g), which dissolved in water forms hydriodic acid HI_(aq). Note the group formula pattern.

 Reaction with Group 1 Alkali Metals Li Na K etc.

• Alkali metals burn very exothermically and vigorously when heated in chlorine to form colourless crystalline ionic salts e.g. NaCl or Na⁺Cl^–. This is a very expensive way to make salt! Its much cheaper to produce it by evaporating sea water!
• e.g. sodium + chlorine ==> sodium chloride
• 2Na_(s) + Cl_2(g) ==> 2NaCl_(s)
• The sodium chloride is soluble in water to give a neutral solution pH 7, universal indicator is green. The salt is a typical ionic compound i.e. a brittle solid with a high melting point. Similarly potassium and bromine form potassium bromide KBr, or lithium and iodine form lithium iodide LiI.  Again note the group formula pattern.

 Reaction with other metals

• If aluminium or iron is heated strongly in a stream of chlorine (or plunge the hot metal into a gas jar of chlorine carefully in a fume cupboard) the solid chloride is formed.
• aluminium + chlorine ==> aluminium chloride_{(white solid)}
  • 2Al_(s) + 3Cl_2(g) ==> 2AlCl_3(s)
• iron + chlorine ==> iron(III) chloride_{(brown solid)}
  • 2Fe_(s) + 3Cl_2(g) ==> 2FeCl_3(s)
• If the iron is repeated with bromine the reaction is less vigorous, with iodine there is little reaction, these also illustrate the halogen reactivity series.