Advanced A level inorganic chemistry exam revision notes for

Part 9: The chemistry of the Group 17 HALOGENS

9.1 An introduction to halogens, group data and trends in the physical and chemical properties of the p-block elements called the halogens - Group 17 fluorine, chlorine, bromine, iodine & astatine

Sub-index for this page on introducing the Group 17 Halogens

(1) Introduction and important trends of the group 7/17 halogen elements

(2) Oxidation states in a variety of compounds and their systematic names

(3) Data table of selected physical and chemical properties - trends discussed and explained

(4) FLUORINE – brief summary of a few points

(5) CHLORINE – brief summary of a few important points

INDEX of all Group 17 Halogens revision notes

Part 9. Group 7/17 The Halogens

 9.1 Introduction and some important trends of the Group 17 Halogen elements

The typical physical properties of the halogens are described and discussed with reference to the data table. Another table describes the patterns in formulae of halogen compounds.

Group 7/17 halogen trends in properties are described, explained and discussed.

There is also a section on how to name halogen compounds.

A brief chemistry of fluorine, bromine, chlorine, iodine and astatine is given which is then elaborated on in the subsequent pages – use the Group 7/17 halogen sub–index

The halogens position in the periodic table and their outer electron configuration is ns²np⁵

(1) The Halogens are typical non–metals and form the 7th Group in the Periodic Table.

The latest modern periodic table denotes it as Group 17, treating the vertical columns of the d–blocks as groups 3 to 12. '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.

• 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).

  • The increase in melting/boiling points down group 7/17 is because the weak electrical intermolecular attractive forces increase with increasing size of atom or molecule.

  • Since they are non–polar molecules, the only intermolecular force that can operate is that which arises from the instantaneous dipole – induced dipole interactions which increase in strength the greater the number of electrons in the molecule.

• They are all coloured non–metallic elements and the colour gets darker down the group.

• They are all poor conductors of heat and electricity – typical of non–metals, there are no free electrons as in metals and no ions either.

• 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 and the outer electrons are increasingly more shielded and less strongly attracted by the nucleus.

• Astatine is highly radioactive, not easily studied for that reason and as far as I know, it has no uses.

• The atoms all have 7 outer electrons (outer electron configuration nS²np⁷ where n = 2 to 6), 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^–. (see halogens bonding page)

  • they form covalent compounds with non–metals and with themselves (see table below). The bonding in the molecule involves single covalent bonds (bond order 1) e.g. hydrogen chloride HCl or H–Cl, double covalent bonds (bond order 2) e.g. chlorine(IV) oxide ClO₂ or O=Cl=O and in the oxyanions the bond order is between 1 and 2 e.g. 1.5 in the delocalised electron systems (see halogens bonding page)

(2) Oxidation states in a variety of compounds and their systematic names

• Oxidation states, examples of formulae and their names

  • A brief summary here, see also redox reactions of halogens and halide ions

  • The oxidation states range from –1, +1, +3, (+4), +5, (+6) and +7 with examples quoted below.

  • The table also exemplifies how to systematically name halogen compounds.

• Apart from 0 in the element, the oxidation state of fluorine is always –1 in compounds – it is the most electronegative element.

• The maximum oxidation state expected is +7, equal to the number of outer valency shell electrons (ns²np⁵).

• The +4 and +6 oxidation states only occur in a few compounds such as some of the halogen oxides.

• 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 KF, hydrogen chloride HCl, sodium chloride NaCl, calcium bromide CaBr₂

    • magnesium iodide MgI₂, etc. for when the oxidation state of the halogen is usually –1 in these simple ionic compounds.

  • The oxides and halides (so called interhalogen compounds) of the halogens are named as halogen(ox.st.) oxide and halogen(ox.st.) halide e.g.

    • bromine(VI) oxide BrO₃,  bromine(III) fluoride BrF₃

  • Other than (i) the 'hydrohalic' acids HX_(aq) (ox. st. –1), the acids derived from halogens in a positive oxidation state are named as (ii) 'halic' (ox.st.) acid and the corresponding anions are named as 'halate'(ox.st.) e.g.

    • (i) hydrochloric acid HCl, chloride ion Cl^-

    • (ii) chloric(V) acid HClO₃, chlorate(I) ion ClO₃^-

• 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 (halogen displacement description).

• The reactivity decreases down the group (explanation of halogen reactivity trend).

• They are all TOXIC elements, in the case of chlorine, in particular, this is put to good use! (for more detail 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!

• How to identify halogens and compounds is in section 9.5 Tests for halide ions – chloride, bromide, iodide

• Miscellaneous points:

  • Preparation of hydrogen chloride and chlorine gases – Method Ex 4

(3) Data table of selected physical and chemical properties

A selection of physical and chemical properties

down group 7/17 ===> halogen trends
property\Z X, name	9F fluorine	17Cl chlorine	35Br bromine	53I iodine	85At astatine (radioactive)
Period	2	3	4	5	6
appearance
	F2 pale yellow gas	Cl2 pale green gas	Br2 dark red liquid, orange–brown gas	I2 dark grey solid, purple vapour	At2 ~black solid, very dark vapour on heating
melting point/oC	–219	–101	–7	114	302
boiling point/oC	–188	–34	59	184	380
density/gcm–3	1.1(l)	1.56(l)	3.12(l)	4.93(s)	approx 7.5(s)
1st IE/kJmol–1	1681	1251	1140	1010	920
covalent atomic radius/pm	64	99	114	133	140
Van der Waal radius pm	155	180	190	195	na
X– ionic radius/pm	136	181	195	216	227
X–X(g) bond enthalpy kJmol–1	158	242	193	151	110
H–X(g) bond enthalpy kJmol–1	562	431	366	299	na
C–X(g) bond enthalpy kJmol–1	484	338	276	238	na
electronegativity	3.98	3.16	2.96	2.66	2.20
electron configuration	2.7	2.8.7	2.8.18.7	2.8.18.18.7	2.8.18.32.18.7
	1s22s22p5	[Ne]3s23p5	[Ar]3d104s24p5	[Kr]4d105s25p5	[Xe]4f145d106s26p5
known oxidation states	–1 only	–1,+1,3,4,5,7	–1,+1,3,4,5,6,7	–1,+1,3,5,7	–1, +1, +3
Electrode potential X2/X–	+2.87V	+1.36V	+1.07V	+0.54V	+0.20V
Electron affinity/kJmol–1	F  –348	Cl  –364	Br  –342	I  –314	At  –270
property\Z X, name	9F fluorine	17Cl chlorine	35Br bromine	53I iodine	85At astatine (radioactive)
******************	*************	***************	*****************	********************	************************

• Some trends have already been discussed in section (1)

• Electronegativity is the power of an atom to attract electron charge from another atom it is covalently bonded to. Some Pauling values of electronegativity are quoted below.

• element	Na	Mg	Al	Mn	Fe	H	Si	P	C	S	I	Br	Cl	N	O	F
  electronegativity	0.9	1.2	1.5	1.5	1.8	2.1	1.8	2.1	2.5	2.5	2.5	2.8	3.0	3.0	3.5	4.0

• Generally speaking electronegativity increases from left to right across a period of the periodic table and decreases down a group of the periodic table.

• As you go down a group the outer electrons are further from the nucleus and increasingly shielded by an extra layer of filled inner quantum levels. So, the net effect is an increasingly weaker pull on the outer electrons by the nucleus. This results in e.g. increasingly lower ionisation energies and increased atomic radii down a group BUT it also weakens the ability of an atom to attract an electron cloud towards it in the context of it sharing bonding electrons when covalently bonded to an atom of a different element.

• Therefore down a group, such as the Group 7/17 Halogens, the electronegativity steadily falls (see table above) and this has major consequences on the bond character in halogen compounds.

• This decrease in ability to attract electrons/electron cloud also means the oxidising power decreases down the group e.g. the electrode potential for F₂/F^- is +2.87V and falls to +0.54V for I₂/I^-. This trend is important in understanding the displacement reactions between the halogen elements and salts of other halogens (halide salts).

Some Group VII (Group 7/17) Halogens trends in bond lengths and bond enthalpies

Some general observations, most of which relate to smaller radii giving shorter stronger bonds:

Halogen molecules X₂:

From bromine to iodine the bond length increases and, except for fluorine, the bond enthalpy decreases as the radius of the halogen atom increases with increasing number of filled inner electron shells. Fluorine is distinctly anomalous with a much lower than expected bond dissociation energy, though the bond length fits the general trend. This is explained by the close proximity of the small fluorine atoms causing repulsion between them due to the closeness of the outer electron orbitals.

Hydrogen halides HX:

From hydrogen fluoride HF_(g) to hydrogen iodide HI_(g), there is clear trend in increasing bond length and decreasing bond enthalpy. One result is the increasing ease of aqueous ionisation from hydrofluoric acid to hydriodic acid so that the HX_(aq) acids become stronger down the group.

HX(aq)  +  H₂O(l)  ===>  H₃O⁺(aq)  +  X^-(aq)

In fact, hydrofluoric acid HF_(aq) is a relatively weak acid but hydrochloric, hydrobromic and hydriodic acids are all very strong. The latter three are so strong in aqueous media you don't really see the difference e.g. from pH readings, but, in non–aqueous polar solvent media, the differences can be clearly measured.

Halogenoalkanes R₃C–X:

Based on polarisation of the bond (C^δ+–X^δ–), you might expect the reactivity order with respect to nucleophiles (electron pair donors) attacking the δ+ carbon bond to be R–F > R–Cl > R–Br > R–I as the electronegativity difference decreases from C–F to C–I. However, it is the decreasing bond enthalpies from C–F to C–I that override this polarisation trend giving the reactivity trend R–I > R–Br > R–Cl > R–F.

See Nucleophilic substitution in halogenoalkanes

(4) FLUORINE – brief summary of a few important points

• The structure of the element:

  • Non–metal existing as covalent diatomic molecule, F₂, with a single bond.

• Physical properties: 

  • Pale yellow gas; mpt –219^oC; bpt –188^oC;  poor conductor of heat/electricity.

• Group, electron configuration (and oxidation states): 

  • Gp7 Halogen; e.c. 2,7  or 1s²2s²2p⁵;  (only –1) e.g. HF, ClF, F₂O (O is +2!)

  • An extremely reactive element and readily combines with almost every other element.

• Reaction of element with oxygen: 

  • None, but oxygen difluoride, F₂O, can be made indirectly.

  • Fluorine is the most reactive element in the periodic table and reacts directly with almost every other element.

  • It is an extremely powerful oxidising agent.

• Reaction of oxide with water:

  • The oxide hydrolyses to form hydrofluoric acid and oxygen (it is powerful enough to oxidise water. This is an anomalous reaction for a Group 7 element due to the high oxidising power of oxygen in the +2 state in F₂O.

    • F₂O_(g) + H₂O_(l) ==> 2HF_(aq) + O_2(g) 

• Reaction of oxide with acids:

  • It readily reacts with water as above.

• Reaction of oxide with bases/alkalis:

  • Presumably fluoride salt is formed and oxygen, as it oxidises water.

• Reaction of element with chlorine: 

  • Can combine directly or indirectly to form ClF, ClF₃, ClF₅ and ClF₇.

  • e.g. Cl_2(g) + F_2(g) ==> 2ClF_(g)

• Reaction of chloride with water:

  • ClF_x + H₂O => ?

• Reaction of element with water:

  • Reacts to form hydrofluoric acid and oxygen. This is an anomalous reaction for a Group 7 element due to the high oxidising power of fluorine.

    • 2F_2(g) + 2H₂O_(l) ==> 4HF_(aq) + O_2(g) 

• Links to other pages on site: 

  • KS4 Science GCSE/IGCSE revision notes on Group 7 The Halogens

  • Advanced Level Chemistry Notes on the rest of the p–block

(5) CHLORINE – brief summary of a few important points

• The structure of the element:

  • Non–metal existing as covalent diatomic molecule, Cl₂, with a single bond.

• Physical properties: 

  • Pale green gas; mpt –101^oC; bpt –34^oC;  poor conductor of heat/electricity.

• Group, electron configuration (and oxidation states): 

  • Gp7 Halogens; e.c. 2,8,7  or 1s²2s²2p⁶3s²3p⁵;  (ranges from –1 to +7) e.g.

  • HCl and NaCl (–1), NaClO and Cl₂O (+1), NaClO₂ (+3), KClO₃ (+5), Cl₂O₇ and HClO₄ (+7).

• Reaction of element with oxygen: 

  • None, but there are several oxides made indirectly.

• Reaction of the oxides with water:

  • Chlorine(I) oxide forms weak chloric(I) acid.

    • Cl₂O_(g) + H₂O_(l) ==> 2HClO_(aq)  

  • Chlorine(V) oxide forms the strong chloric(VII) acid.

    • Cl₂O_7(l) + H₂O_(l) ==> 2HClO_4(aq) 

• Reaction of oxide with acids:

  • None, only acidic in nature.

• Reaction of oxide with bases/alkalis:

  • chlorine(I) oxide forms sodium chlorate(I) with sodium hydroxide,

    • Cl₂O_(g) + 2NaOH_(aq) ==> 2NaClO_(aq) + H₂O_(l)

    • ionic equation: Cl₂O_(g) + 2OH^–_(aq) ==> 2ClO^–_(aq) + H₂O_(l)

  • and chlorine(VII) oxide will dissolve to form sodium chlorate(VII)

    • Cl₂O_7(l) + 2NaOH_(aq) ==> 2NaClO_4(aq) + H₂O_(l)

    • ionic equation: Cl₂O_7(l) + 2OH^–_(aq) ==> 2ClO₄^–_(aq) + H₂O_(l)

• Reaction of element with water:

  • Slightly reacts to form an acid mixture of chloric(I) acid and hydrochloric acid, but the position of the equilibrium is very much on the left.

    • Cl_2(g) + H₂O_(l) HClO_(aq) + HCl_(aq) 

    • or more accurately and ionically ...

    • Cl_2(g) + 2H₂O_(l)  HClO_(aq) + H₃O⁺_(aq) + Cl^–_(aq) 

• Other comments:

  • –

• Links to other pages on site: 

  • KS4 Science GCSE/IGCSE revision notes on Group 7 The Halogens

  • Advanced Level Chemistry Notes on the rest of the p–block