Group 7 HALOGEN elements of the Periodic Table

See also Salt - sodium chloride - extraction - uses of halogens

(c) doc bDoc Brown's Chemistry KS4 science GCSE/IGCSE/O level Chemistry Revision Notes

Group 7 of the Periodic Table – The Halogens (non-metals)

(c) doc b (c) doc b (c) doc b (c) doc b (c) doc b

The halogens – fluorine, chlorine, bromine, iodine, astatine

Revising Physical Properties, Chemical Reactions & Uses of the Group VII Halogen elements and their compounds

particularly, salt (sodium chloride) and the many products derived from it in the chlor-alkali industry

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. The balanced molecular equations and ionic equations of the reactions of halogens, explaining the reactivity trend of the Group VII halogen elements, the uses of the halogens, uses of halide salts and halogen organochlorine compounds. These revision notes on the halogens should prove useful for the new AQA, Edexcel and OCR GCSE (9–1) chemistry science courses.

THINKING AHEAD: From a working knowledge of the position of the Group 7 Halogen elements in the periodic table you should be able to predict the number of outer electrons of Group 7 Halogen elements, possible compound formulae of the Group 7 Halogens, reactions and symbol equations for Group 7 Halogens and the probable reactivity of  a halogen in group 7 from its position in the periodic table and the physical properties of elements low down in the group like astatine.

Advanced Level Students GCE Advanced AS A2 IB Level Chemistry Group 7/17 Halogen Notes

BUT reading this page reminds you of what you theoretically leaned from GCSE/IGCSE/O Level courses on halogens!

So this page can act as a primer for the study of the halogens chlorine, bromine iodine etc.

(c) doc b 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?

(c) doc b

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. The halogens are next to the last element in any period from period 2 onwards.

The non-metallic halogens have seven outer electrons, in any period from period 2 onwards. This outer electron similarity of the halogens makes them behave in a chemically similar (e.g. similar formulae) and in a particularly reactive way and is a modern pre-requisite of a set of elements belonging to the same group. BUT their similarity in physical properties and chemical reactions fits in well with Mendeleev's original conception of a group classification.

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

  • 2.7fluorine ==> 2.8.7(c) doc bchlorine ==> down group 7 all have  seven outer electrons

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

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

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

  • Halogens all exist as diatomic molecules (F2, Cl2, Br2 etc. see diagram above!)

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

  • The halogens are all coloured non–metallic elements and the colour gets darker down Group 7 (see data table) and diagram above. At room temperature ...

    • Fluorine is a pale yellow gas, the most reactive non-metallic element known

    • Chlorine a dense pale green gas, highly reactive and very toxic (used in WWI).

    • Bromine is a dark red liquid that is easily vapourised to an orange–brown vapour, toxic and still pretty reactive

    • Iodine is a very dark grey crystalline solid that on heating sublimes to give a brilliant purple vapour, not as reactive as the others

    • Astatine is a very dark, almost black, solid, that gives an equally dense dark coloured vapour on heating, the least reactive

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

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

  • The density of the halogens increases down Group 7 halogens.

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

  • Health and safety issues and the Group 7 Halogens hazard hazard hazard

    • All of the Group 7 halogens are harmful and irritants, particularly if the vapours are breathed in, and chlorine and iodine are toxic.

    • All the vapours irritate the respiratory system and can cause lung damage, so any experiments should be done in a fume cupboard.

    • Fluorine is not only the most reactive halogen, its the most reactive non-metal and has to be handled with extreme care.

    • Fluorine attacks glass forming silicon fluoride, so any apparatus containing it must be specially coated glass or special metal alloys, both of which must have extremely inert surfaces.

    • Chlorine was used as a gas attack agent in the First World War, causing blindness, lung damage and death.

    • Liquid bromine is very corrosive and must not come into contact with the skin and breathing in bromine vapour is not a good idea!


Chemical features, similarities, physical property and reactivity trends

Elements - non-metals on the on the far right-hand side of the periodic table, and apart from noble gases, quite readily gain an electron into their outer shell, giving them a high reactivity in forming negative ions or a covalent bond. The group 7 halogens have 7 outer electrons and only need to gain one electron to form a stable negative ion, or, share one electron to make a covalent bond, thus making them the most reactive non-metals of the periodic table.

  • The halogen 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.

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

      • Halogen atoms gain one negative electron (reduction) to be stable and this gives a surplus electric charge of –1.

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

    • Halogens form ionic compounds with metals e.g. sodium chloride Na+Cl. (ionic bonding revision notes page)

    • BUT halogens also form covalent compounds with non–metals and with themselves (see below).

  • Note on naming halogen compounds:

    • When halogens combine with other elements in simple ionic 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. and collectively known as the halides.

  • 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.

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

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

  • (c) doc bAstatine 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 advanced data)

Symbol and Name of halogen

Atomic Number of halogen Electron arrangement of halogen State and colour of halogen  at room temperature and pressure, colour of vapour when heated Melting point of halogens Boiling point of halogens atom radius of halogens nm (nanometre), pm (picometre)

F Fluorine

9 2.7 (c) doc bpale yellow gas –219oC, 54K –188oC 85K 0.064, 64

Cl Chlorine

17 2.8.7 (c) doc bpale green gas –101oC, 172K  –34oC, 239K 0.099, 99
Br Bromine 35 (c) doc bdark red liquid, readily gives off a brown vapour –7oC, 266K 59oC, 332K 0.114, 114

I Iodine

53 (c) doc b dark (~black) crumbly solid, purple vapour 114oC, 387K 184oC, 457K 0.133, 133
At Astatine 85

(c) doc b(c) doc bdark solid, dark vapour – highly radioactive! 302oC, 575K 380oC 653K 0.140, 140



Proton number of group 7 halogens


All group 7 halogens have seven electrons in the outer shell.


Down group 7 the halogen's colour gets darker.



The melting points of group 7 halogens increase down the group.


The boiling points of group 7  increase down the group.


The atomic radii of group 7 halogens increases down the group.


Note: For atomic radii: 1nm = 10–9m,  1pm = 10–12m,  nm x 1000 = pm,  nm = pm/1000

Atomic radii always increase down a group with increase in atomic number because extra electron shells are successively added.



The Bonding in Group 7 Halogen Compounds

When halogens combine with other metals, ionic compounds are formed e.g. sodium chloride, magnesium chloride and aluminium fluoride, where the ionic bonding results from the attraction of the positive metal ion and the negative halide ion, formed by electron transfer.

ONE (c) doc b atom combines with ONE (c) doc b atom to form (c) doc b(c) doc b

ONE (c) doc b atom combines with TWO (c) doc b atoms to form (c) doc b(c) doc b

ONE (c) doc b atom combines with THREE (c) doc batoms to form (c) doc b(c) doc b


When the non–metallic halogens combine with other non–metals the result is a covalent bond in a covalent compound molecule, formed by electron sharing e.g.

(c) doc b and (c) doc b combine to form (c) doc b hydrogen chloride by electron sharing

(c) doc ba chlorine atoms in a molecule of chloromethane, a C–Cl covalent bond

(c) doc b a bromine atom in a molecule of bromoethane, a C–Br covalent bond

The Reactivity Order and Displacement Reactions

A more reactive halogen will displace a less reactive halogen from a solution of its compounds.

Such experiments provide an experimental method for establishing the reactivity order down the group of halogens.

An experiment to demonstrate the order of reactivity fro the Group 7 Halogens.

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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

i.e. the more reactive halogen displaces a less reactive halogen

The word and symbol equations for the 1 – 3 DISPLACEMENT REACTIONS on the diagram are given below.

The equations are given below with and without state symbols followed by the ionic equations for the halogen displacement reactions.

NOTE: From a suitable safe supply, you can bubble chlorine gas through a bromide or iodide solution, but this will be a teacher demonstration, its much better that a class of students can do the experiments for themselves quite safely using chlorine water!

By learning how to write and balance the symbol equation for the reaction between ANY halogen and a halogen salt or halide ion of lesser reactive halogen, all you have to do is get the symbol of another group 7 halogen from your periodic table and deduce the equation i.e. swap Br for an I etc.

1. chlorine + potassium bromide ==> potassium chloride + bromine

Cl2 + 2KBr ==> 2KCl + Br2

Cl2(aq) + 2KBr(aq) ==> 2KCl(aq) + Br2(aq)

2. chlorine + potassium iodide ==> potassium chloride + iodine

Cl2 + 2KI ==> 2KCl + I2

Cl2(aq) + 2KI(aq) ==> 2KCl(aq) + I2(aq)

3. bromine + potassium iodide ==> potassium bromide + iodine

Br2 + 2KI ==> 2KBr + I2


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.  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.

In terms of oxidation (electron loss) and reduction (electron gain) in this 'redox' reaction ...

Chlorine molecules gain electrons and are reduced to chloride ions.

The bromide ion is oxidised because it loses an electron in forming bromine molecules.


The other two possible reaction equations involving (2) chlorine + iodide and (3) bromine + iodide, are similar to the example above and are shown below.

2.  Cl2(aq) + 2I(aq) ==> 2Cl(aq) + I2(aq)

Chlorine molecules gain electrons and are reduced to chloride ions.

The iodide ion is oxidised because it loses an electron in forming iodine molecules.


3.  Br2(aq) + 2I(aq) ==> 2Br(aq) + I2(aq)

Bromine molecules gain electrons and are reduced to bromide ions.

The iodide ion is oxidised because it loses an electron in forming iodine molecules.


Predicting the chemistry of astatine At

Astatine is highly radioactive element and dangerous to work with in the laboratory and any prior knowledge or expectation would allow safer working. However, by following the group 7 halogen element patterns and trends you can make good predictions as to how astatine will behave physically and chemically.

For example, from the reactivity trend, you would expect astatine to be less reactive than iodine (above it) because the halogens get less reactive down the group i.e. F > Cl > Br > I > At

Therefore you would expect e.g. chlorine, bromine or iodine to displace astatine from its salts, therefore, these three predicted halogen displacement reaction are, in word equations, in symbol equations and ionic redox equations as follows ...

1. chlorine + potassium astatide ==> potassium chloride + astatine

Cl2(aq) + 2KAt(aq) ==> 2KCl(aq) + At2(aq)

Cl2(aq) + 2At(aq) ==> 2Cl(aq) + At2(aq)

2. bromine + potassium astatide ==> potassium bromide + astatine

Br2(aq) + 2KAt(aq) ==> 2KBr(aq) + At2(aq)

Br2(aq) + 2At(aq) ==> 2Br(aq) + At2(aq)

3. iodine + potassium astatide ==> potassium iodide + astatine

I2(aq) + 2KAt(aq) ==> 2KI(aq) + At2(aq)

I2(aq) + 2At(aq) ==> 2I(aq) + At2(aq)

Note that you can predict the formulae of the astatine molecule and potassium astatide because astatine is in the same group as Cl, Br and I with the same number of outer electrons, same valency, therefore the formulae will fit into the same pattern e.g. the At2 diatomic element molecule, the salt potassium astatide KI. If you follow that, say to yourself,  well done me and Mendeleev !!!


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?

(c) doc b  F [2.7] + e ==> (c) doc bF [2.8]

(c) doc b Cl [2.8.7] + e ==> (c) doc bCl [2.8.8]

Br [] + e ==> Br []

I [] + e ==> I []

  • The electronic structure of the halogens is the basis for explaining their high reactivity and reactivity order.

  • 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) because they are one electron short of forming a stable full outer shell.

    • Or, as in this case of the halogen displacement reactions explained here, they gain an electron to form a stable singly charged negative ion with a Noble Gas electron arrangement (diagrams above for fluorine and bromine and the electron arrangement changes for bromine and iodine).

  • 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)

    • This process of forming a negative ion by the process of electron gain is an example of reduction.

  • 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 inner electron shell, AND this extra shell of electrons also has a 'shielding' effect, so the combined effect is that the outer electrons are progressively less strongly held (less strongly attracted) as you go down the group.

    • e.g. 2.7fluorine ==> 2.8.7(c) doc bchlorine ==> down group 7

    • This means the further (and more shielded) the negative electrons are from the positive nucleus the less strongly they are held, but this also means the smaller the atom the more strongly electrons are attracted, and as the atomic radius decreases up the group (normal group trend), the smaller the atom, the more strongly electrons are attracted, and the halogens want to attract an electron to form a stable singly charged negative ion with a noble gas electron arrangement.

    • So up the group 7 halogens the outer electrons are closer and closer to the nucleus and less shielded, so the halogen atom can attract an electron more strongly, meaning they become reactive up the group.

    • The smaller atoms attract 'incoming' electrons more strongly to form a halide ion (or shared to form a covalent bond for that matter).

  • 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

QUALITATIVE TESTS FOR HALIDE IONS – the negative ions (anions) formed from the halogens

To the suspected halide ion solution add a little dil. nitric acid and a few drops of silver nitrate solution.

Depending on the halide ion you get a different coloured silver halide precipitate, summarised below.

halide ion Colour of precipitate with silver nitrate Ionic equation to show precipitate formation
chloride Cl white precipitate of AgCl silver chloride (slowly darkens when exposed to light - see use in photography) Ag+(aq) + Cl(aq) ==> AgCl(s)
bromide Br cream precipitate of AgBr silver bromide Ag+(aq) + Br(aq) ==> AgBr(s)
Iodide I yellow precipitate of  AgI silver iodide  Ag+(aq) + I(aq) ==> AgI(s)

Note: A simple chemical test for chlorine:

Chlorine is the only common gas that bleaches litmus paper. If you use blue litmus it turns pink first and is then bleached white.

More details on these and other tests


Other Reactions of the Halogens

Fluorine forms fluorides, chlorine forms chlorides and iodine forms iodides,

and these compounds maybe ionic with metals or covalent with other non-metals

for other important industrial reactions see Salt - sodium chloride - extraction - uses of halogens

Reaction with hydrogen H2

  • (c) doc bHalogens readily combine with hydrogen to form the hydrogen halides which are colourless gaseous covalent molecules.

  • e.g. hydrogen + chlorine ==> hydrogen chloride

  • H2(g) + Cl2(g) ==> 2HCl(g)

  • A combination of two non-metals gives a low melting/boiling covalent compound.

  • BUT, watch out for complications, all the hydrogen halides e.g. hydrogen fluoride HF, hydrogen chloride HCl, hydrogen bromide HBr and hydrogen iodide HI, ALL dissolve in water to give strongly acid ionic solutions!

    • The hydrogen halides dissolve in water to form very strong acids with solutions of pH0 to pH1

    • e.g. hydrogen chloride forms hydrochloric acid in water HCl(aq) or H+Cl(aq) because the hydrogen halides (HCl, HBr and HI) are all fully ionised in aqueous solution even though the original hydrogen halides were covalent!

    • Theory reminder - an acid is a substance that forms H+ ions in water.

    • See also Uses of Halogens and halogen compounds including sodium chloride and chlorine

  • 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, so you can predict how astatine will behave and the formula of its compounds.

  • Fluorine, chlorine and iodine readily react with other non-metals like phosphorus and sulfur - no need to know the formula of these covalent compounds.



Most halogens readily react with metals, especially on heating, to form ionic compounds

The charge on the halide ion formed is -1

Reaction with Group 1 Alkali Metals Li Na K etc.

  • (c) doc bAlkali metals burn very exothermically and vigorously when heated in chlorine to form colourless crystalline ionic salts e.g. NaCl or Na+Cl.

    • You can do this by setting up a glass tube containing a lump of alkali metal, in a fume cupboard, heat the metal and pass chlorine gas through the glass tube and over the hot alkali metal - very exothermic, may well burn with a bright flame.

    • A cloud of white/colourless crystals of the metal chloride will settle out somewhere downstream!

    • 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) + 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 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.

  • e.g. if you heat alkali metals with bromine (in a fume cupboard) you get white/colourless crystals of the bromide salts, which of course, are also ionic compounds, formed by electron transfer.

    • e.g. potassium + bromine ==> potassium bromide

      • 2K + Br2 ==> 2KBr

      • The reactions with bromine are less vigorous and exothermic compared to chlorine, and reactions between alkali metals and iodine are even less reactive i.e. you still observe the general group 7 reactive trend of 'less reactive down the group'.

  • Complete ionic bonding details revision notes on another page.

  • For safe ways of reacting metals with chlorine see the apparatus illustrated in the next section.


Reaction of halogens 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.

    • The apparatus for the preparation of aluminium chlorise (c) doc b

      • aluminium + chlorine ==> aluminium chloride(white solid)

        • 2Al(s) + 3Cl2(g) ==> 2AlCl3(s)

  • Similarly you can heat iron wire (wool better) in a stream of chlorine with similar apparatus.

    • prep FeCl3

      • iron + chlorine ==> iron(III) chloride(brown solid)

        • 2Fe(s) + 3Cl2(g) ==> 2FeCl3(s)

  • If the experiment with iron is repeated with bromine the reaction is less vigorous, and with iodine there is little reaction

    • These iron plus halogen reactions also illustrate the halogen reactivity series.

    • I have managed to demonstrate the reaction of iron with bromine or iodine by heating them together in a wide pyrex test tube, in a fume cupboard and you can clearly see the difference in reaction of chlorine, bromine and iodine towards the same metal iron (fair test and all that !).

    • For more details see method of making aluminium chloride and iron(III) chloride


See also Salt - sodium chloride - extraction - uses of halogens


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

BUT reading this page reminds you of what you theoretically leaned from GCSE/IGCSE/O Level courses on halogens!

So this page can act as a primer for the study of the halogens chlorine, bromine iodine etc.



OTHER RELATED WEBPAGES separate web pages


PLEASE NOTE that these LINKS are for Advanced Level (GCE A level) Students ONLY

ADVANCED LEVEL INORGANIC CHEMISTRY Part 9 Group 7/17 Halogens sub–index: 9.1 Introduction, trends & Group 7/17 data * 9.2 Halogen displacement reaction and reactivity trend  * 9.3 Reactions of halogens with other elements * 9.4 Reaction between halide salts and conc. sulfuric acid * 9.5 Tests for halogens and halide ions * 9.6 Extraction of halogens from natural sources * 9.7 Uses of halogens & compounds * 9.8 Oxidation & Reduction – more on redox reactions of halogens & halide ions * 9.9 Volumetric analysis – titrations involving halogens or halide ions * 9.10 Ozone, CFC's and halogen organic chemistry links * 9.11 Chemical bonding in halogen compounds * 9.12 Miscellaneous aspects of halogen chemistry

See also Salt - sodium chloride - extraction - uses of halogens

keywords equations: astatine * bromine * chemical characteristics * chlorine * data on the elements * displacement reaction *  explaining reactivity trend * fluorine * hydrogen halides * iodine * naming halogen compounds * physical characteristics * reaction with metals * 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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