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Advanced Level Inorganic Chemistry Periodic Table Revision Notes Part 9. Group 7/17 The Halogens

9.1 An Introduction to halogens, group data and trends in the properties of the halogens

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 below.

 

PLEASE NOTE KS4 Science GCSE/IGCSE/O Level GROUP 7 HALOGENS NOTES are on a separate webpage

INORGANIC 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

Advanced Level Inorganic Chemistry Periodic Table Index * Part 1 Periodic Table history * Part 2 Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies * Part 3 Period 1 survey H to He * Part 4 Period 2 survey Li to Ne * Part 5 Period 3 survey Na to Ar * Part 6 Period 4 survey K to Kr and important trends down a group * Part 7 s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals * Part 8  p–block Groups 3/13 to 0/18 * Part 9 Group 7/17 The Halogens * Part 10 3d block elements & Transition Metal Series * Part 11 Group & Series data & periodicity plots * All 11 Parts have their own sub–indexes near the top of the pages


Part 9. Group 7/17 The Halogens

 9.1 Introduction and some important trends

 

Pd s block d blocks and f blocks of metallic elements p block elements
Gp1 Gp2 Gp3/13 Gp4/14 Gp5/15 Gp6/16 Group7/17 Gp0/18
1

1H

2He
2 3Li 4Be The modern Periodic Table of Elements

ZSymbol, z = atomic or proton number

highlighting position of Group 7/17 Halogens

5B 6C 7N 8O 9F

fluorine

10Ne
3 11Na 12Mg 13Al 14Si 15P 16S 17Cl

chlorine

18Ar
4 19K 20Ca 21Sc 22Ti 23V 24Cr 25Mn 26Fe 27Co 28Ni 29Cu 30Zn 31Ga 32Ge 33As 34Se 35Br

bromine

36Kr
5 37Rb 38Sr 39Y 40Zr 41Nb 42Mo 43Tc 44Ru 45Rh 46Pd 47Ag 48Cd 49In 50Sn 51Sb 52Te 53I

iodine

54Xe
6 55Cs 56Ba 57-71 72Hf 73Ta 74W 75Re 76Os 77Ir 78Pt 79Au 80Hg 81Tl 82Pb 83Bi 84Po 85At

astatine

86Rn
7 87Fr 88Ra 89-103 104Rf 105Db 106Sg 107Bh 108Hs 109Mt 110Ds 111Rg 112Cn 113Uut 114Fl 115Uup 116Lv 117Uus

ununseptium

118Uuo

outer electrons ns2np5

Pd s block d blocks and f blocks of metallic elements p block elements
Gp1 Gp2 Gp3/13 Gp4/14 Gp5/15 Gp6/16 Group 7/Group 17 Gp0/18
1

1H 1s1

2He 1s2
2 3Li [He]2s1 4Be [He]2s2 Electronic structure of selected elements of the periodic table

ZSymbol, Z = atomic/proton number = total electrons in neutral atom

elec. config. abbreviations: [He] = 1s2 [Ne] = 1s22s22p6

[Ar] = 1s22s22p63s23p6     [Kr] = 1s22s22p63s23p63d104s24p6

5B [He]2s22p1 6C [He]2s22p2 7N [He]2s22p3 8O [He]2s22p4 9F

[He]2s22p5

10Ne [He]2s22p6
3 11Na [Ne]3s1 12Mg [Ne]3s2 13Al [Ne]3s23p1 14Si [Ne]3s23p2 15P [Ne]3s23p3 16S [Ne]3s23p4 17Cl

[Ne]3s23p5

18Ar [Ne]3s23p6
4 19K [Ar]4s1 20Ca [Ar]4s2 21Sc [Ar] 3d14s2 22Ti [Ar] 3d24s2 23V [Ar] 3d34s2 24Cr [Ar] 3d54s1 25Mn [Ar] 3d54s2 26Fe [Ar] 3d64s2 27Co [Ar] 3d74s2 28Ni [Ar] 3d84s2 29Cu [Ar] 3d104s1 30Zn [Ar] 3d104s2 31Ga [Ar] 3d104s24p1 32Ge [Ar] 3d104s24p2 33As [Ar] 3d104s24p3 34Se [Ar] 3d104s24p4 35Br

[Ar] 3d104s24p5

36Kr [Ar] 3d104s24p6
5 37Rb [Kr]5s1 38Sr [Kr]5s2 39Y [Kr] 4d15s2 40Zr [Kr] 4d25s2 41Nb [Kr] 4d45s1 42Mo [Kr] 4d55s1 43Tc [Kr] 4d55s2 44Ru [Kr] 4d75s1 45Rh [Kr] 4d85s1 46Pd [Kr] 4d10 47Ag [Kr] 4d105s1 48Cd [Kr] 4d105s2 49In [Kr] 4d105s25p1 50Sn [Kr] 4d105s25p2 51Sb [Kr] 4d105s25p3 52Te [Kr] 4d105s25p4 53I

[Kr] 4d105s25p5

54Xe [Kr] 4d105s25p6
6 55Cs [Xe]6s1 56Ba [Xe]6s2 4f–block and 5d–block in period 6 including Lanthanide Series 81Tl [Xe] 4f145d106s26p1 82Pb [Xe] 4f145d106s26p2 83Bi [Xe] 4f145d106s26p3 84Po [Xe] 5d106s26p4 85At

[Xe] 4f145d106s26p5

86Rn [Xe] 4f145d106s26p6
7 87Fr [Rn]7s1 88Ra [Rn]7s2 5f–block & 6d–block including Actinide Series of Metals in period 7 113Uut [Rn] 5f146d107s27p1 114Fl [Rn] 5f146d107s27p2 115Uup [Rn] 5f146d107s27p3 116Lv [Rn] 5f146d107s27p4 117Uus

[Rn] 5f146d107s27p5

118Uuo [Rn] 5f146d107s27p6
  **************************  

 

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), this 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.

  • 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.

  • 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.

  • (c) doc b 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)

  • Advanced Inorganic Chemistry Page Index and LinksOxidation states

    • 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.

oxidation state of halogen in the compound –1 +1 +3 +4 +5 +6 +7
examples of compounds or ions NaCl sodium chloride

CaF2 calcium fluoride

AlBr3 aluminium bromide

F2O oxygen(II) fluoride

NaOCl sodium chlorate(I)

Cl2O chlorine(I) oxide

HClO chloric(I) acid

ClF chlorine(I) fluoride

KClO3 potassium chlorate(III)

BrO2 bromate(III) ion

ClF3 chlorine(III) fluoride

ClO2  chlorine(IV) oxide

BrO2  bromine(IV) oxide

ClO3 chlorate(V) ion

I2O5 iodine(V) oxide

HIO3 iodic(V) acid

BrO3 bromine(VI) oxide HClO4 chloric(VII) acid

IO4 iodate(VII) ion

Cl2O7 chlorine(VII) oxide

  ****************** ************ **************** ************* ************** ************ ************
  • 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.

  • Advanced Inorganic Chemistry Page Index and LinksNote 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. for when the oxidation state is –1.

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

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

  • 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).

  • (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!

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

  • Advanced Inorganic Chemistry Page Index and LinksMiscellaneous points:

down group 7/17 ===>
property\Z X, name 9F fluorine 17Cl chlorine 35Br bromine 53I iodine 85At astatine (radioactive)
Period 2 3 4 5 6
appearance (c) doc b (c) doc b (c) doc b (c) doc b (c) doc b
pale yellow gas pale green gas dark red liquid, orange–brown gas dark grey solid, purple vapour 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 energy kJmol–1 158 242 193 151 110
H–X(g)bond energy kJmol–1 562 431 366 299 na
C–X(g)bond energy 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 0.20
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)
**************************** *************** ******************* ********************* ******************** ***************************
  • 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 bond character in halogen compounds.

Advanced Inorganic Chemistry Page Index and Links


FLUORINE – brief summary of a few points

  • The structure of the element:

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

  • Physical properties

    • Pale yellow gas; mpt –219oC; bpt –188oC;  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.

  • 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) + O2(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. Cl2(g) + F2(g) ==> 2ClF(g)

  • Reaction of chloride with water:

    • ClFx + 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.

      • 2F2(g) + 2H2O(l) ==> 4HF(aq) + O2(g) 

  • Links to other pages on site

Advanced Inorganic Chemistry Page Index and Links


CHLORINE – brief summary of a few points

  • The structure of the element:

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

  • Physical properties

    • Pale green gas; mpt –101oC; bpt –34oC;  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.

      • Cl2O7(l) + H2O(l) ==> 2HClO4(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)

      • Cl2O7(l) + 2NaOH(aq) ==> 2NaClO4(aq) + H2O(l)

      • ionic equation: Cl2O7(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.

      • Cl2(g) + H2O(l) HClO(aq) + HCl(aq) 

      • or more accurately and ionically ...

      • Cl2(g) + 2H2O(l)  HClO(aq) + H3O+(aq) + Cl(aq) 

  • Other comments:

  • Links to other pages on site

Advanced Inorganic Chemistry Page Index and Links


PLEASE NOTE KS4 Science GCSE/IGCSE/O Level GROUP 7 HALOGENS NOTES are on a separate webpage

WHAT NEXT?

INORGANIC 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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