The Halogens – Group 7

THE HALOGENS – GROUP 7

The halogens are all in group 7 on the right of the periodic table. This group consists of elements like Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), Astatine (At).

The Halogens are typical non-metals and form the 7th Group in the Periodic Table ‘Halogens’ means ‘salt formers’ and the most common compound is sodium chloride which is found from natural evaporation as huge deposits of ‘rock salt’ or the even more abundant as ‘sea salt’ in the seas and oceans.

 

Physical properties

• Typical non-metals with relatively low melting points and boiling points.
• The melting points and boiling increase steadily down the group (so the change in state at room temperature from gas => liquid => solid), this is because the inter molecular attractive forces increase with increasing size of atom or molecule.
• They are all coloured non-metallic elements.
• The colour of the halogen gets darker down the group.
• They are all poor conductors of heat and electricity – typical of non-metals.
• When solid they are brittle and crumbly e.g. iodine.
• The size of the atom gets bigger as more inner electron shells are filled going down from one period to another.

 

Chemical properties

The atoms all have 7 outer electrons, this outer electron similarity, as with any Group in the Periodic Table, makes them have very similar chemical properties eg

• They form singly charged negative ions e.g. chloride Cl- because they are one electron short of a noble gas electron structure.

They gain one negative electron (reduction) to be stable and this gives a surplus electric charge of -1.

These ions are called the halide ions, two others you will encounter are called the bromide Br- and iodide I- ions.

• They form ionic compounds with metals e.g. sodium chloride Na+Cl-. They form covalent compounds with non-metals and with themselves.

 

The bonding in the molecule involves single covalent bonds e.g. hydrogen chloride HCl or H-Cl

• The elements all exist as X2 or X-X, diatomic molecules where X represents the halogen atom.
• The reactivity decreases down the group.
• They are all TOXIC elements.
• Astatine is very radioactive, so difficult to study but its properties can be predicted using the principles of the Periodic Table and the Halogen Group trends!

 

Group 7 (Halogens)

These have their outermost electron orbits with electrons one less than that required for complete filling.

These elements are gaseous in nature and have valence -1, they borrow electrons to stabilize their electronic configuration.They are all diatomic covalently bonded molecules. Diatomic means that each molecule contains two atoms.

The formulae are F₂, Cl₂, Br₂, I₂, (see structure of chlorine).

All of the halogens will either.

1) Gain one electron from a metal to form an ionic bond, or

2) Share one electron with a non-metal to form a covalent bond

 

Reactions with Metals.

The halogens will gain one electron to form an ionic bond with metals. The halogens will react with

1) Group 1 Metals

• Alkali metals burn very exothermically and vigorously when heated in chlorine to form colourless crystalline ionic salts eg NaCl or Na⁺Cl^–.

This is a very expensive way to make salt! It’s much cheaper to produce it by evaporating sea water!

E.g. sodium + chlorine ==> sodium chloride

2Na(s) + Cl₂(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 ie 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.

 

1) Transition metals

Iron + bromine iron (III) bromide.

2Fe(s) + 3Br₂(l) 2FeBr₃(s)

Iron + chlorine iron(III) chloride.

2Fe(s) + 3Cl₂(g) 2FeCl₃(s) All of the compounds with metals are ionic salts which form a giant structure.

They are called metal halides because they are formed from a metal and a halogen.

 

Testing halide ions.

Halide ions undergo a series of unique reactions that allow an unknown solid or aqueous sample to be tested for the presence of chloride, bromide or iodide ions.

Aqueous silver ions react with halide ions to produce individually coloured precipitates. These precipitates have different solubilities in ammonia solution and so further differentiation can be achieved, see the table below,

Aqueous lead(II) ions react with halide ions to produce a different set of halide salts, see the table below,

 

Summary of the Reactivity Trend of Halogens

• All Group VII metals have 7 valence electrons (7 electrons in the highest energy level)
• Atomic radius increases down the Group as successive ‘electron shells’ (energy levels) are filled
• Electro negativity (the relative tendency shown by a bonded atom to attract electrons to itself) decreases down the group as the elements become more metallic in nature.

(Typically, metals have low electro negativity, little ability to attract electrons, while non metals have high electro negativity, greater ability to attract electrons).

 

The reactivity of Group VII elements is related to the element’s ability to attract electrons, so the greater the electro negativity, the more reactive the Halogen.

So, chemical reactivity of Group VII elements decreases down the Group, from the most reactive (Fluorine) to the least reactive (Iodine).

• Down the Group, first ionization energy (the energy required to remove 1 electron from the gaseous atom) decreases. As the atomic radius increases and the electron is further from the nucleus it is less attracted to the nucleus (electron is said to be ‘shielded’)
• Melting point and boiling point increase down the Group as the elements become more metallic in nature
• There is a gradation in color going down the group, the elements become darker in colour as they become more metallic in nature.

Similarly there is a gradation in physical appearance at STP, from gas to liquid to solid, as the elements become more metallic in nature.

 

Uses of Halogens

1. Chlorine by itself is used as bleach and in the manufacture of sodium chlorate, which can be used as bleach and a herbicide.
2. Water purification also relies on chlorine to kill bacteria in the water, after the impure water has passed through various filtration stages.
3. Chlorine is also used in the production of chlorofluorocarbons, commonly called CFC’s, used in the past as refrigerant gases and propellants for aerosol cans.

Both these uses have now been banned by international law, in the developed world at least.

4. The problem these chemicals cause is that when they reach the high atmosphere the molecules break apart to release chlorine atoms.

 

These chlorine atoms then react with molecules of ozone, O³, turning them into molecules of oxygen. One chlorine atom can destroys thousands of molecules of ozone.

Ozone traps UV light from the sun, preventing it hitting the surface of the Earth, and with a depletion of ozone more UV light gets through, which increases occurrences of skin cancer in humans.

5. Fluorine is used as fluoride salts in toothpaste or added to domestic water supplies to strengthen teeth enamel helping to minimise tooth decay. (eg potassium fluoride).
6. Bromine and iodine are both used in ‘halogen’ car headlamps.
7. Iodine is used in hospitals in the mild antiseptic solution ‘tincture of iodine’.

 

See also:

TREND IN ATOMIC RADIUS

REACTIVITY TREND OF ALKALI METALS

THE PERIODIC TABLE

ISOTOPES

THE STRUCTURE OF THE ATOM AND THE PERIODIC TABLE