Group 7 HALOGEN elements of the Periodic Table

Doc Brown's Chemistry KS4 science GCSE/IGCSE/O level Chemistry Revision Notes

See also Salt - sodium chloride - extraction - uses of halogens and halogen compounds

Group 7 of the Periodic Table – The physical and chemical properties of the halogens (non-metals)

Sub-index GCSE notes on the Group 7 Halogens

1. Where are the Group 7 Halogens in the Periodic Table?

2. Electronic structure and the reactivity of Group 7 Halogens (non-metals)

3. A general introduction to the Halogens

4. Chemical features, similarities, physical property data, and reactivity trends

5. Covalent and ionic bonding in Group 7 halogen compounds

6. The reactivity order and halogen displacement reactions

7. Oxidation–Reduction theory of halogen displacement reactions

8. Explaining the reactivity trend of the Group 7 halogens

9. Predicting the chemistry of astatine At

10. Qualitative tests for the halogens and halide ions

11a. Reactions of the halogens with hydrogen, production of acids

11b. Reactions of the halogens with metals, formation of both ionic and covalent compounds

GCSE foundation tier easier QUIZ on the Group 7 Halogens

GCSE higher tier–harder QUIZ on the Group 7 Halogens

For advanced A Level student see Advanced A 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.

KEYWORDS - a sort of sub-index!

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

1. 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!

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. Group 7 elements are on the far right of the periodic table with 7 outer electrons (1 short of a noble gas structure) and so you would expect them to be very reactive non-metals and form singly charged negative ions.  It is the similarity in electron structure (7 electrons in the outer shell) that makes the chemistry of group 7 halogen non-metals the same - group 7 chemistry!

2. Electronic structure and reactivity of Group 7 Halogens (non-metals)

In the context of their position in the Periodic Table

On reaction non–metals readily form negative ions in compounds by gaining electrons e.g. 

chlorine ==> chloride: Cl₂ + 2e^– ===> 2Cl^– (more simply Cl + e^– ===> Cl^– typical of Group 7 Halogens)

oxygen ===> oxide: O₂ + 4e^– ==> 2O^2– (more simply O + 2e^– ===> O^2– typical of Group 6 elements)

These are typical electron changes when non-metallic elements in groups 6 and 7 react.

The negative ions are formed directly from the non-metals like halogen atoms.

Atoms usually react to give an electron arrangement with a full outer shell by losing, gaining or sharing electrons.

Non-metallic elements on the on the far right-hand side of the periodic table, (apart from the very noble gases which already have a stable full outer shell), quite readily gain electrons into their outer shell, giving them a high reactivity in forming negative ions.

The outer electrons of non-metals tend to be more strongly held than the outer electrons of metals and this is very much the case for group 7 halogens which are the elements the furthest on the right of the periodic table (bar the stable noble gases).

Therefore, the group 7 halogens like fluorine, chlorine and bromine tend to be the most reactive non-metallic elements.

For non-metals, it usually takes too much energy to remove to many electrons to give a stable positive ion electron arrangement, but its much easier for a non-metal, like those in group 6 or 7, to gain 2 or 1 electrons to give an electronically stable negative ion with a full outer shell of electrons like a noble gas.

Group 6 and 7 elements also readily share the outer electrons of other non-metals to form covalent bonds

e.g. H₂O and H₂S from group 6 (O, S)  and  for the group 7 halogens like chlorine, HCl and CCl₄.

Non-metals like group 7 halogens do NOT normally form positive ions. You would have to remove 7 electrons from a chlorine atom to make the ion Cl⁷⁺, and this requires far too much energy that any chemical reaction could deliver!

The group 7 halogens require to gain or share the least electrons to form an ion or molecule in which the halogen atom has a very stable noble gas electron arrangement. This requires the least energy, so the group 7 halogens tend to be the most reactive non-metals on the right-hand side of the periodic table.

These points and explanations are elaborated on by looking at the chemical reactions of halogens further down the page.

3. A general 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, similar reactions) 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 ==> 2.8.7chlorine ==> down group 7 elements 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?

• General properties and trends down the Group 7 Halogens with increase in atomic number and relative atomic mass

• 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 of pairs of atoms)

  • The atoms or molecules of the halogens get larger down the group because from one halogen down to the next, an extra shell of electrons is added e.g. in terms of electron arrangement: F 2.7, Cl 2.8.7, Br 2.8.18.7 etc.

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

  • 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 increasingly weak electrical intermolecular attractive forces with increasing size of atom or molecule.

    • Generally speaking the more electrons in the molecule the greater the intermolecular attractive force between molecules as they increase in size.

• 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).

    • A simple test for chlorine: The gas turns damp blue litmus red and then bleaches it white.

  • 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

    • A simple test for iodine: It turns starch solution a blue-black colour.

  • 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

  • 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!

4. 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).

    • 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 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.

  • 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 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	pale yellow gas, extremely reactive and toxic	–219oC, 54K	–188oC 85K	0.064, 64
Cl Chlorine	17	2.8.7	pale green gas, toxic	–101oC, 172K	–34oC, 239K	0.099, 99
Br Bromine	35	2.8.18.7	dark dense red liquid, readily gives off a brown vapour, toxic	–7oC, 266K	59oC, 332K	0.114, 114
I Iodine	53	2.8.18.18.7	dark (~black) crumbly solid, purple vapour	114oC, 387K	184oC, 457K	0.133, 133
At Astatine	85	2.8.18.32.18.7	dark solid, dark vapour – highly radioactive!	302oC, 575K	380oC 653K	0.140, 140
The GROUP 7 HALOGENS   ***************	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.

5. Covalent and ionic 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  atom combines with ONE atom to form

ONE  atom combines with TWO atoms to form

ONE  atom combines with THREE atoms to form

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.

 and combine to form hydrogen chloride by electron sharing

a chlorine atoms in a molecule of chloromethane, a C–Cl covalent bond

a bromine atom in a molecule of bromoethane, a C–Br covalent bond

You can use these trends to make predictions about the physical and chemical properties of group 7 halogen elements

6. The Reactivity Order and Halogen 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.

7. Oxidation–Reduction theory of halogen displacement reactions

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.

In terms of oxidation (electron loss) and reduction (electron gain) in these 'redox' reactions ...

Reduction: Chlorine molecules gain electrons and are reduced to chloride ions (gain of one electron per chlorine atom).

Oxidation: The bromide ion is oxidised because it loses an electron in forming bromine atoms ==> molecules (loss of one electron per bromide ion).

 

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

Reduction: Chlorine molecules gain electrons and are reduced to chloride ions (gain of one electron per chlorine atom).

Oxidation: The iodide ion is oxidised because it loses an electron in forming iodine atoms ==> molecules (loss of one electron per iodide ion).

 

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

Reduction: Bromine molecules gain electrons and are reduced to bromide ions (gain of one electron per bromine atom).

Oxidation: The iodide ion is oxidised because it loses an electron in forming iodine atoms ==> molecules (loss of one electron per iodide ion).

8. 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?

Pd

metals

metals

non-metal group

Gp1

Gp2

Gp3

Gp4

Gp5

Gp6

Gp7

Gp0

1

He

2

Li

Be

a short section of the periodic table with group 7 electron arrangements

B

C

N

O

9F 2.7

Ne

3

Na

Mg

Al

Si

P

S

17Cl 2.8.7

Ar

4

K

Ca

Sc

Ti

V

Cr

Mn

Fe

Co

Ni

Cu

Zn

Ga

Ge

As

Se

35Br 2.8.18.7

Kr

  F [2.7]   +   e^–   ===>  F^– [2.8]^–   (for the formation of the fluoride ion)

Cl [2.8.7]   +   e^–   ===>   Cl^– [2.8.8]^–    (for the formation of the chloride ion)

Br [2.8.18.7] + e^– ===> Br^– [2.8.18.8]^–   (for the formation of the bromide ion)

I [2.8.18.18.7] + e^– ===> I^– [2.8.18.18.8]^–   (for the formation of the iodide ion)

• 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 chlorine 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).

    • Strictly speaking the correct half-equation is:  Cl₂ + 2e^–  ===> 2Cl^–   (see the redox equations above)

  • 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 has a 'shielding' effect on the outer shell electrons which are less strongly attracted the further they are from the nucleus.

  • 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 ==> 2.8.7chlorine ==> 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 the outer electrons are attracted, hence the more reactive the halogen.

  • As the atomic radius decreases up the group, the smaller the atom, the more strongly electrons are attracted, the more the halogen atom wants 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.

  • Final comment: Fluorine is the most reactive element of the periodic table in the sense, as far as I know, it reacts with nearly every other element in the periodic table, often violently and very exothermically to form very strong ionic or covalent bonds with the other elements.

  • Fluorine also reacts with the vast majority of compounds of other elements too!

  • The only elements fluorine doesn't react with are the noble gases helium, neon and argon, surprise! surprise!.

9. Predicting the chemistry of astatine At - one of the most important ideas in using the periodic table

(You won't find much on astatine in your textbooks! and even less on tennessine Ts, the 6th halogen)

Astatine, element 85, is a 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.

You would expect astatine to consist of diatomic molecules (At₂) have a higher melting point and boiling point than iodine.

You expect it to form a singly charged negative ion, the astatide ion At^-, by an astatine atom with 7 electrons in the outer shell, gaining one electron

You would expect it to form ionic compounds with metals e.g. sodium astatide Na⁺At^-  or calcium astatide CaAt₂ etc.

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

Cl_2(aq) + 2KAt_(aq) ==> 2KCl_(aq) + At_2(aq)

Cl_2(aq) + 2At^–_(aq) ==> 2Cl^–_(aq) + At_2(aq)

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

Br_2(aq) + 2KAt_(aq) ==> 2KBr_(aq) + At_2(aq)

Br_2(aq) + 2At^–_(aq) ==> 2Br^–_(aq) + At_2(aq)

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

I_2(aq) + 2KAt_(aq) ==> 2KI_(aq) + At_2(aq)

I_2(aq) + 2At^–_(aq) ==> 2I^–_(aq) + At_2(aq)

You can also predict that astatine will form hydrogen astatide (HAt) that is very soluble in water to give a strong acid solution of pH 0-1, just like hydrochloric acid (HCl) etc.

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 ionic salt potassium astatide KI.

 If you have followed this lot, say to yourself, well done me and Mendeleev !!!

Postscript: A few atoms of the 7th halogen element tennessine (Ts) have been made.

They are highly unstable and radioactive, but if you could make sufficient of them, would Ts behave as a halogen and fit in with the group patterns?

The answer is probably yes and you could substitute Ts for At in any of the above formulae and equations.

It would be the least reactive and the highest melting very dark solid, assuming you predict from the group 7 halogen trends.

In other words you would expect it to follow the trend set by the previous five halogens.

i.e. melting point trend:  Ts > At > I > Br > Cl > F

and reactivity trend:  F > Cl > Br > I > At > Ts

10. 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.

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 (chlorine is acidic in water) and is then the litmus paper bleached white.

AND, in biology you test for starch using a dilute iodine solution - you get a deep blue-black colour.

Therefore a simple test for iodine in a solution is to add a drop of starch and see if you get the blue-black colour.

More details on these and other tests

11a. 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

• Halogens readily combine with hydrogen to form the hydrogen halides which are colourless gaseous covalent molecules.

  • Complete notes on covalent bonding.

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

• H_2(g) + Cl_2(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.

• When halogens are combined with non-metallic elements, covalent compounds are formed.

11b. REACTION of HALOGENS with METALS

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

The charge on the halide ion formed is -1

Reaction of halogens 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^–.

  • 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.

  • The sodium loses one electron to form a positive ion and the electron is transferred to a chlorine atom to form a chloride ion, their mutual attraction constitutes the ionic bond. (ionic bonding)

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

  • You can also do the experiment by heating a small amount of the alkali metal in a deflagrating spoon and plunging into a gas jar of previously prepared chlorine when the metal will burn quite vigorously.

  • 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.

• 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.

All the alkali metals react with all of the halogens to produce white crystalline solids of the ionic compound.

e.g sodium reacts with fluorine gas to give sodium fluoride.

2Na(s)  +  F₂(g)  ===>  2NaF(s)

Since the charge on the group 1 metal ions is +1, and the charge on halide ions is -1, its easy to predict the formula of any ionic compound formed between an alkali metal and a halogen i.e. a 1 : 1 ratio.

Gp1\7	F	Cl	Br	I
Li	LiF	LiCl	LiBr	LiI
Na	NaF	NaCl	NaBr	NaI
K	KF	KCl	KBr	KI
Rb	Rb	RbCl	RbBr	RbI
Cs	CsF	CsCl	CsBr	CsI

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.

  • 

    • aluminium + chlorine ==> aluminium chloride_{(white solid)}

      • 2Al_(s) + 3Cl_2(g) ==> 2AlCl_3(s)

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

  • 

    • iron + chlorine ==> iron(III) chloride_{(brown solid)}

      • 2Fe_(s) + 3Cl_2(g) ==> 2FeCl_3(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

• These two chlorides vapourise quite easily on heating and their bonding is closer to covalent, than ionic.

• When halogens are combined with metallic elements, ionic compounds are usually formed, but not always.

WHERE–WHAT NEXT?

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

Preparation of hydrogen chloride & chlorine gases is described on gas preparation page Method Ex 4

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 your GCSE/IGCSE/O Level courses on halogens!

So this page can act as a primer for the study of the group 7/17 halogens chlorine, bromine iodine etc.

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

• GCSE/IGCSE electrochemistry notes with more on electrolysis

• Summary of the chemistry of the Periodic Table

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

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