Group 7 of the Periodic Table – The Halogens

Revising Physical Properties, Chemical Reactions & Uses

The physical properties of the Group 7 halogens – fluorine, chlorine, bromine, iodine and astatine are described and detailed notes on the chemical displacement reactions of chlorine, bromine and iodine with balanced molecular and ionic equations, explaining the reactivity trend of the Group VII halogen elements, the uses of the halogens, uses of halide salts and halogen organochlorine compounds

PLEASE NOTE A Level Students GCE Advanced AS A2 IB Level Chemistry Group 7 Halogen Notes

KEYWORDS: astatine * bleach * bromine * chemical characteristics * chlorine * data on the elements * displacement reaction * electrolysis of NaCl * explaining reactivity trend * fluorine * hydrochloric acid * hydrogen halides * iodine * naming halogen compounds * physical characteristics * PVC * reaction of sodium hydroxide and chlorine * reaction with metals * reaction with hydrogen * silver halide photography * uses of chlorine * uses of fluorine, bromine and iodine * uses of hydrogen * uses of sodium chloride * uses of sodium hydroxide

Where are the Group 7 Halogens in the Periodic Table?

The Group VII Halogens form the next to the last vertical column on the right of the Periodic Table, where you find most of the non–metallic elements. Therefore the Halogen is the next to the last element on the period from period 2 onwards. At the bottom of Group 7 is the radioactive halogen astatine (At) which is not shown.

Note: Using 0 to denote the Group number of Noble Gases is very historic now since compounds of xenon known exhibiting a valency of 8. Because of the horizontal series of elements e.g. like the Sc to Zn block (10 elements), Groups 3 to 0 can also be numbered as Groups 13 to 18 to fit in with the actual number of vertical columns of elements. This can make things confusing, but there it is, classification is still in progress!

Introduction to the Halogens (see also halogens data table below)

The Halogens are typical non–metals and form the 7th Group in the Periodic Table (the vertical pink column above). '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 'sea salt' in the seas and oceans.

Physical features and important trends down the Group with increasing atomic number (proton number)

• What are the group trends in melting point, boiling point, reactivity, size of atom (atomic radius), density as you go down the group 7 as the atomic/proton number increases?

• Its helpful to compare the trends with the halogen elements information in the halogens data table

• They are typical non–metals with relatively low melting points and boiling points.

• The melting points and boiling of the Halogens increase steadily down Group 7 (so the change in state at room temperature from gas ==> liquid ==> solid)

  • Why do the melting points and boiling points of Group 7 Halogens increase with atomic number?, i.e. increase down the group

  • This increase in melting/boiling points down Group 7 is due to the increasing weak electrical intermolecular attractive forces with increasing size of atom or molecule.

    • Generally the more electrons in the molecule the greater the attractive force between molecules.

• They are all coloured non–metallic elements and the colour gets darker down Group 7 (see data table).

• They are all poor conductors of heat and electricity – typical of non–metals.

• When solid they are brittle and crumbly e.g. iodine.

• The density increases down Group 7.

• The size of the atom gets bigger down Group 7 as more inner electron shells are filled going down from one period to another.

Chemical features, similarities, and physical property and reactivity trends

• 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 e.g.

  • 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^–. (ionic bonding revision notes page)

  • they form covalent compounds with non–metals and with themselves (see below). The bonding in the molecule involves single covalent bonds e.g. hydrogen chloride HCl or H–Cl. (covalent bonding revision notes page)

• Note on naming halogen compounds:

  • When combined with other elements in simple compounds the name of the halogen element changes slightly from ...ine to ...ide.

  • Fluorine forms a fluoride (ion F^–), chlorine forms a chloride (ion Cl^–), bromine a bromide (ion Br^–) and iodine an iodide (ion I^–).

  • The other element at the start of the compound name e.g. hydrogen, sodium, potassium, magnesium, calcium, etc. remains unchanged.

  • So typical halogen compound names are,  potassium fluoride, hydrogen chloride, sodium chloride, calcium bromide, magnesium iodide etc.

• The elements all exist as X₂ or X–X, diatomic molecules where X represents the halogen atom.

• A more reactive halogen can displace a less reactive halogen from its salts.

  • A halogen displacement description

• The reactivity of halogens decreases down Group 7 with increase in atomic number.

  • The explanation of the Group 7 Halogen reactivity trend

• they are all TOXIC elements, which has its advantages in some situations! (See uses of Halogens)

• 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!

• Details of how to identify halogens and their compounds are on the Chemical Tests page (use the alphabetical list at the top of this other page) – Tests for halide ions – chloride, bromide, iodide

DATA Selected Properties of the Group 7 Halogens (more Group 7 halogens AS–A2 data)
Symbol and Name	Atomic Number	Electron arrangement	State and colour at room temperature and pressure, colour of vapour when heated	Melting point	Boiling point	atom radius pm
F Fluorine	9	2.7	pale yellow gas	–219oC, 54K	–188oC 85K	64
Cl Chlorine	17	2.8.7	pale green gas	–101oC, 172K	–34oC, 239K	99
Br Bromine	35	2.8.18.7	dark red liquid, readily gives off a brown vapour	–7oC, 266K	59oC, 332K	114
I Iodine	53	2.8.18.18.7	dark (~black) crumbly solid, purple vapour	114oC, 387K	184oC, 457K	133
At Astatine	85	2.8.18.32.18.7	black solid, dark vapour – highly radioactive!	302oC 575K	380oC 653K	140

The Reactivity Order and Displacement Reactions

What happens when a solution of a halogen like chlorine, bromine or iodine is added to a salt solution of another halogen i.e. adding a halogen to a sodium/potassium chloride, sodium/potassium bromide or sodium/potassium iodide solution? Observations! What do you see when a halogen displacement reaction happens? A few drops of chlorine water, bromine water and iodine water are added in turn to aqueous solutions of the salts  potassium chloride (KCl), potassium bromide (KBr) and potassium iodide (KI). Three combinations produce a reaction (and three don't!). You can get 'simple' observations from the diagrams! A darkening effect compared to a water blank confirms a displacement reaction has happened. Chlorine displaces bromine from potassium bromide and iodine from potassium iodide.  Bromine only displaces iodine from potassium iodide and the least reactive iodine cannot displace chlorine or bromine from their salts. On the basis that the most reactive element displaces a least reactive element the reactivity order must be: chlorine > bromine > iodine The word and symbol equations for the 1 – 3 DISPLACEMENT REACTIONS on the diagram are given below.

1. chlorine + potassium bromide ==> potassium chloride + bromine Cl2(aq) + 2KBr(aq) ==> 2KCl(aq) + Br2(aq) 2. chlorine + potassium iodide ==> potassium chloride + iodine Cl2(aq) + 2KI(aq) ==> 2KCl(aq) + I2(aq) 3. bromine + potassium iodide ==> potassium bromide + iodine Br2(aq) + 2KI(aq) ==> 2KBr(aq) + I2(aq) The halogen molecule is the electron acceptor (the oxidising agent) and is reduced by electron gain to form a halide ion The halide ion is the electron donor (the reducing agent) and is oxidised by electron loss to form a halogen molecule chlorine molecule + bromide ion ==> chloride ion + bromine molecule ionically the redox equations are written ... 1.  Cl2(aq) + 2Br–(aq) ==> 2Cl–(aq) + Br2(aq) because the potassium ion, K+, is a spectator ion, that is, it does not take part in the reaction. The other two possible reaction equations involving (ii) chlorine + iodide and (iii) bromine + iodide, are similar to the example above. 2.  Cl2(aq) + 2I–(aq) ==> 2Cl–(aq) + I2(aq) 3.  Br2(aq) + 2I–(aq) ==> 2Br–(aq) + I2(aq)

The colours of the observations as shown on the diagram have been a bit simplified. The more concentrated the potassium halide salt solutions and the more halogen solution you add e.g. chlorine/bromine/iodine water, the deeper the colours formed. e.g. the grid below matches the diagram BUT with more subtle and wider ranging observations that you are likely to see in reality. The bold observations 1., 2. and 3.indicate a major colour change i.e. a displacement reaction has happened and match the diagram and equations above.

Halogen added	KCl solution	KBr solution	KI solution	BLANK of water
chlorine Cl2	VERY pale green solution	1. orange–reddish brown solution	2. brown solution–black precipitate	VERY pale green solution
bromine Br2	orange–reddish brown solution	orange–reddish brown solution	3. brown solution–black precipitate	orange–reddish brown solution
iodine I2	dark brown solution	dark brown solution	dark brown solution	dark brown solution

Higher GCSE level Oxidation–Reduction Theory

The halogen molecule is the electron acceptor (the oxidising agent) and is reduced by electron gain to form a halide ion

The halide ion is the electron donor (the reducing agent) and is oxidised by electron loss to form a halogen molecule

chlorine molecule + bromide ion ==> chloride ion + bromine molecule

ionically the redox equations are written ...

1.  Cl_2(aq) + 2Br^–_(aq) ==> 2Cl^–_(aq) + Br_2(aq)

because the potassium ion, K+, is a spectator ion, that is, it does not take part in the reaction. The other two possible reaction equations involving (ii) chlorine + iodide and (iii) bromine + iodide, are similar to the example above.

2.  Cl_2(aq) + 2I^–_(aq) ==> 2Cl^–_(aq) + I_2(aq)

3.  Br_2(aq) + 2I^–_(aq) ==> 2Br^–_(aq) + I_2(aq)

Explaining the Reactivity Trend of the Group 7 Halogen

Why are Group 7 Halogen non–metallic elements so reactive?

Why do halogens get less reactive down the group with increase in atomic/proton number?

How do we explain the reactivity trend of the group 7 halogens?

  F [2.7] + e^– ==> F^– [2.8]^–

Cl [2.8.7] + e^– ==> Cl^– [2.8.8]^–

Br [2.8.18.7] + e^– ==> Br^– [2.8.18.8]^–

I [2.8.18.18.7] + e^– ==> I^– [2.8.18.18.8]^–

• The halogen atoms are only one electron short of pseudo Noble Gas electron arrangement, which are particularly stable. Therefore, in a chemical reaction, halogens try to complete the outer octet of electrons by forming a single covalent bond (sharing a pair of electrons) or, as in this case of halogen displacement reactions explained here, they gain an electron to form a stable singly charged negative ion with a Noble Gas electron arrangement.

• In halogen displacement reactions, when a halogen atom reacts, it gains an electron to form a singly negative charged ion

  • e.g. Cl + e^–  ==> Cl^– which has a stable noble gas electron structure like argon. (i.e. 2.8.7 ==> 2.8.8)

• As you go down the group with increase in atomic number from one Group 7 halogen element down to the next .. F => Cl => Br => I ...

  • the atomic radius gets bigger due to an extra filled electron shell,

  • the outer electrons are further and further from the nucleus and are also shielded by the extra full electron shell of negative electron charge,

  • therefore the outer electrons are less and less strongly attracted by the positive nucleus as would be any 'incoming' electrons to form a halide ion (or shared to form a covalent bond).

• SO, this combination of factors means to attract an 8th outer electron is more and more difficult as you go down the group, so the element is less reactive as you go down the group, i.e. less 'energetically' able to form the X^– halide ion with increase in atomic number.

The preparation of hydrogen chloride and chlorine gases is described on the gas preparation page Method Ex 4

Other Reactions of the Halogens

note: fluorine forms fluorides, chlorine forms chlorides and iodine forms iodides

Reaction with hydrogen H₂

• Halogens readily combine with hydrogen to form the hydrogen halides which are colourless gaseous covalent molecules. Complete covalent bonding details revision notes on another page.

• 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.

• Complete ionic bonding details revision notes on another page.

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.

  • (see diagram of method of making aluminium chloride and iron(III) chloride) in 4. methods of making salts)

• 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.

Qualitative Analysis – Tests for halide ions

• These are all described on the chemical Identification pages

The Uses of Chlorine, the brine electrolysis products and other halogens and their compounds

 

Sodium Chloride NaCl	sodium chloride is an important raw material found as 'rock salt' or in seawater. Salt is used directly for food flavouring and food preservation. It is also used for de–icing roads because it lowers the freezing point of water (related to general effect of an impurity lowering the melting point of a substance). BUT sodium chloride is mainly converted into other products by electrolysis (see above for the production of chlorine, hydrogen and sodium hydroxide by electrolysis of brine (aqueous sodium chloride solution)).
CHLORINE Cl2	All the Halogens are potentially harmful substances and chlorine in particular is highly toxic. It is dangerous to ingest halogens or breathe in any halogen gas or vapour. Chlorine is used to kill bacteria and so sterilise water for domestic supply or in in swimming pools. The sodium hydroxide and chlorine can be chemically combined at room temperature to make the bleach, sodium chlorate(I) NaClO. This is used in some domestic cleaning agents, it chemically 'scours' and chemically 'kills' germs! sodium hydroxide + chlorine ==> sodium chloride + sodium chlorate(I) + water 2NaOH(aq) + Cl2(aq) ==> NaCl(aq) + NaClO(aq) + H2O(l) More details for Advanced Level Chemistry Students Organic phenolic chlorine compounds are used antiseptics and  disinfectants like 'Dettol' or 'TCP' Organic chlorine compounds are used as pesticides, including the now mainly banned DDT. Chlorine is used in making CFC refrigerant gases/liquids but their production and use are being reduced. They break down in the upper atmosphere and the chlorine atoms catalyse the destruction of ozone O3 which absorbs harmful uv radiation. PVC: Chlorine (from electrolysis of NaCl) and ethene (from cracking oil fractions) are used to make a chemical called chloro(ethene), which used to be called vinyl chloride, this is then converted into the plastic–polymer poly(chloroethene) or PVC, because it is shorthand for the old name polyvinylchloride! (equation below) Poly(chloroethene), old names PVC, from chloroethene (vinyl chloride) is much tougher than poly(ethene) and very hard wearing with good heat stability. so it is used for covering electrical wiring and plugs. It is also replacing metals for use as gas and water drain pipes and has found a use as artificial leather and readily dyed to bright colours! (old names : polyvinyl chloride, shortened to PVC) Liquid organic chlorine compounds are used as dry cleaning or de–greasing solvents. PVC is very tough hard wearing useful plastic and a good electrical insulator and is used for water piping, window frames, part of electrical fittings e.g. plug covers etc. Chlorine is used in the manufacture of potassium chlorate weed killer, KClO3.
	HCl(g => aq) As described above, some of the hydrogen and chlorine from the electrolysis of sodium chloride solution are combined to form hydrogen chloride gas. H2(g) + Cl2(g) ==> 2HCl(g)  This gas is dissolved in water to manufacture hydrochloric acid. HCl(g) + aq ==> HCl(aq) or ==> H+(aq) + Cl–(aq)  This is a very important acid used in the chemical industry to make chloride salts.
silver salts Ag+X–	Silver chloride (AgCl), silver bromide (AgBr) and silver iodide (AgI) are all sensitive to light ('photosensitive'), and all three are used in the production of various types of photographic film used to detect visible light and beta and gamma radiation from radioactive materials. Each silver halide salt has a different sensitivity to light. When radiation hits the film the silver ions in the salt are reduced by electron gain to silver (Ag+ + e– ==> Ag, the halide ion is oxidised to the halogen molecule 2X– ==> X2 + 2e– ). AgI is the least sensitive and used in X–ray radiography, AgCl is the most sensitive and used in 'fast' film for cameras, and AgBr is used in most standard photographic films – but much of their use is being superceded by digital cameras!
The other halogens FLUORINE F2 BROMINE Br2 IODINE I2	Fluorine is used as fluoride salts in toothpaste or added to domestic water supplies to strengthen teeth enamel helping to minimise tooth decay. (e.g. potassium fluoride). Fluorine is used in the manufacture of the tough non–stick plastic PTFE coating of cooking pans. Fluorine is used in manufacture of aerosol propellants and refrigerant gases. Apart from its silver salt use in photography, bromine is used to manufacture organic pesticides and fungicides because of their poisonous nature Organic bromine compounds are used as flame inhibitor chemicals (flame retardants) for plastic products to reduce their flammability and as petrol additives to reduce the build–up of lead in car engines (a use decreasing as 'green' unleaded fuels are becoming more popular). Bromine and iodine are both used in 'halogen' car headlamps. Iodine is used in hospitals in the mild antiseptic solution 'tincture of iodine'.
HYDROGEN H2	Hydrogen is used in the manufacture of ammonia (for fertilisers), margarine (by adding hydrogen to unsaturated fats) and hydrochloric acid. It isn't a halogen, but it is made from the electrolysis of salt solution.
SODIUM HYDROXIDE NaOH	Sodium hydroxide is used in the manufacture of soaps, detergents, paper, ceramics and to make soluble salts of organic acids with low solubility in water (e.g. soluble Aspirin). It isn't a halogen compound, but it is made from the electrolysis of salt solution. The sodium hydroxide and chlorine can be chemically combined at room temperature to make the bleach, sodium chlorate(I) NaClO. This is used in some domestic cleaning agents, it chemically 'scours' and chemically 'kills' germs! sodium hydroxide + chlorine ==> sodium chloride + sodium chlorate(I) + water 2NaOH(aq) + Cl2(aq) ==> NaCl(aq) + NaClO(aq) + H2O(l) More details for Advanced Level Chemistry Students

GCE AS/A2/IB Advanced AS A2 IB Level Chemistry Notes on The Halogens

WHERE–WHAT NEXT?

• OTHER RELATED WEBPAGES separate web pages

  • KS4 Science/GCSE/IGCSE Group 7 Halogens task/worksheet

  • GCSE/IGCSE foundation tier–easier quiz on the Group 7 Halogens

  • GCSE/IGCSE higher tier–harder quiz on the Group 7 Halogens

  • A GCSE/IGCSE Group 7 "Halogens" task sheet worksheet * (answers)

  • GCSE/IGCSE word–fill work sheet on the Halogens

  •  The Halogens (matching pair quiz on their appearance)

  • GCSE/IGCSE word–fill worksheet on the electrolysis of sodium chloride solution, products, uses

  • Tests for halide ions – chloride, bromide, iodide

  • GCSE/IGCSE electrochemistry notes with more on electrolysis

  • Summary of the Periodic Table

  PLEASE NOTE A Level Students GCE Advanced AS A2 IB Level Chemistry Group 7 Halogen Notes

keywords equations: astatine * bleach * bromine * chemical characteristics * chlorine * data on the elements * displacement reaction * electrolysis of NaCl * explaining reactivity trend * fluorine * hydrochloric acid * hydrogen halides * iodine * naming halogen compounds * physical characteristics * PVC * reaction of sodium hydroxide and chlorine * reaction with metals * reaction with hydrogen * silver halide photography * uses of chlorine * uses of fluorine, bromine and iodine * uses of hydrogen * uses of sodium chloride * uses of sodium hydroxide * Cl2(aq) + 2KBr(aq) ==> 2KCl(aq) + Br2(aq) * Cl2(aq) + 2KI(aq) ==> 2KCl(aq) + I2(aq) * Br2(aq) + 2KI(aq) ==> 2KBr(aq) + I2(aq) * Cl2(aq) + 2Br–(aq) ==> 2Cl–(aq) + Br2(aq) * Cl2(aq) + 2I–(aq) ==> 2Cl–(aq) + I2(aq) * Br2(aq) + 2I–(aq) ==> 2Br–(aq) + I2(aq) * Cl2(aq) + 2Br–(aq) ==> 2Cl–(aq) + Br2(aq) * Cl2(aq) + 2I–(aq) ==> 2Cl–(aq) + I2(aq) * Br2(aq) + 2I–(aq) ==> 2Br–(aq) + I2(aq) * H2(g) + Cl2(g) ==> 2HCl(g) * 2Na(s) + Cl2(g) ==> 2NaCl(s) * 2Al(s) + 3Cl2(g) ==> 2AlCl3(s) * 2Fe(s) + 3Cl2(g) ==> 2FeCl3(s) * 2NaCl (aq) + 2H2O (l) + elec. energy ==> 2NaOH (aq) + H2 (g) + Cl2 (g) * 2NaOH(aq) + Cl2(aq) ==> NaCl(aq) + NaClO(aq) + H2O(l) * H2(g) + Cl2(g) ==> 2HCl(g) * Cl2 + 2KBr ==> 2KCl + Br2 * Cl2 + 2KI ==> 2KCl + I2 * Br2 + 2KI ==> 2KBr + I2 * Cl2 + 2Br– ==> 2Cl– + Br2 * Cl2 + 2I– ==> 2Cl– + I2 * Br2 + 2I– ==> 2Br– + I2 * Cl2 + 2Br– ==> 2Cl– + Br2 * Cl2 + 2I– ==> 2Cl– + I2 * Br2 + 2I– ==> 2Br– + I2 * H2 + Cl2 ==> 2HCl * 2Na + Cl2 ==> 2NaCl * 2Al + 3Cl2 ==> 2AlCl3 * 2Fe + 3Cl2 ==> 2FeCl3 * 2NaCl + 2H2O + elec. energy ==> 2NaOH + H2 + Cl2  * 2NaOH + Cl2 ==> NaCl + NaClO + H2O * H2 + Cl2 ==> 2HCl *

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