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 level 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
level foundation tier easier
QUIZ on the Group 7 Halogens
GCSE level higher tier–harder QUIZ on the Group 7 Halogens
For advanced A Level student see
Advanced 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.
Where next? Associated pages
KEYWORDS - a sort of sub-index for these Group 7 halogen notes!
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
Doc
Brown's Chemistry KS4 science GCSE/IGCSE/O level Chemistry Revision Notes
Keywords: 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 halogens – fluorine, chlorine,
bromine, iodine, astatine, their physical properties, their
chemical reactions and reactivity. 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.
1. Where are the Group 7 Halogens
in the Periodic Table?
Pd |
metals |
Part of the modern Periodic Table
Pd = period,
Gp = group |
metals => non–metals |
Gp1 |
Gp2 |
Gp3 |
Gp4 |
Gp5 |
Gp6 |
Gp7 |
Gp0 |
1 |
1H Note that
hydrogen does not readily fit into any group |
2He |
2 |
3Li |
4Be |
atomic number
Chemical Symbol eg 4Be |
5B |
6C |
7N |
8O |
9F |
10Ne |
3 |
11Na |
12Mg |
13Al |
14Si |
15P |
16S |
17Cl |
18Ar |
4 |
19K |
20Ca |
21Sc |
22Ti |
23V |
24Cr |
25Mn |
26Fe |
27Co |
28Ni |
29Cu |
30Zn |
31Ga |
32Ge |
33As |
34Se |
35Br |
36Kr |
5 |
37Rb |
38Sr |
39Y |
40Zr |
41Nb |
42Mo |
43Tc |
44Ru |
45Rh |
46Pd |
47Ag |
48Cd |
49In |
50Sn |
51Sb |
52Te |
53I |
54Xe |
6 |
55Cs |
56Ba |
Transition Metals |
81Tl |
82Pb |
83Bi |
84Po |
85At |
86Rn |
|
87Fr |
88Ra |
|
|
|
|
|
|
|
Group 7 The Halogen
*
Reactive Metals of Groups 1 and 2
*
Transition Metals
Post-transition metals - diagonally down and across Groups 3 to
6
The
zig-zag 'line' roughly separates metals from non-metals (except
'metallic' Te)
Non-metals - diagonally down and across Groups 3 to 7 (except
Te)
The very
unreactive Group 0 noble gas non-metals |
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 n on–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– ==> 2O2– (more simply O + 2e–
===> O2–
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.
Pd |
metals |
metals |
non-metal group |
Gp1 |
Gp2 |
Gp3 |
Gp4 |
Gp5 |
Gp6 |
Gp7 |
Gp8 |
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 |
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.
-
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
-
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–.
-
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)
-
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 |
Halogen Atomic Number |
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, pm |
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, reactive and 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 and 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
nm = nanometre and pm = picometre
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.
An experiment to
demonstrate the order of reactivity fro the Group 7 Halogens.
|
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 word equations are given below with the symbol equations with
and without state symbols followed
by the ionic equations for the halogen displacement reactions,
including the theory of oxidation and reduction.
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! |
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 |
|
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
Br2(aq)
+ 2KI(aq) ===> 2KBr(aq) + I2(aq)
Oxidation and reduction analysis involving electron loss
or gain is described in the next section.
|
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. 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 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.
Cl2(aq)
+ 2I–(aq) ==> 2Cl–(aq) + I2(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.
Br2(aq) + 2I–(aq) ==> 2Br–(aq) + I2(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, all showing 7 electrons in the outer
shell |
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,
similarly for the other halide ions ...
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).
-
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
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)
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.
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 (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.
-
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).
-
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!
-
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.
-
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.
Pd |
metals |
Part of the modern Periodic Table
Pd = period,
Gp = group |
metals => non–metals |
Gp1 |
Gp2 |
Gp3 |
Gp4 |
Gp5 |
Gp6 |
Gp7 |
Gp0 |
1 |
1H Note
that H does not readily fit into any group
|
2He |
2 |
3Li |
4Be |
atomic number
Chemical Symbol eg 4Be
Group 1 Alkali Metals and Group 7 Halogens |
5B |
6C |
7N |
8O |
9F |
10Ne |
3 |
11Na |
12Mg |
13Al |
14Si |
15P |
16S |
17Cl |
18Ar |
4 |
19K |
20Ca |
21Sc |
22Ti |
23V |
24Cr |
25Mn |
26Fe |
27Co |
28Ni |
29Cu |
30Zn |
31Ga |
32Ge |
33As |
34Se |
35Br |
36Kr |
5 |
37Rb |
38Sr |
39Y |
40Zr |
41Nb |
42Mo |
43Tc |
44Ru |
45Rh |
46Pd |
47Ag |
48Cd |
49In |
50Sn |
51Sb |
52Te |
53I |
54Xe |
6 |
55Cs |
56Ba |
Transition Metals |
81Tl |
82Pb |
83Bi |
84Po |
85At |
86Rn |
So, you can deduce the
following table of all the possible ionic compounds formed from
the Group 1 Alkali Metals and the non-metal Group 7 Halogen
elements.
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.
-
-
Similarly you can heat
iron wire (wool better) in a stream of chlorine with similar
apparatus.
-
-
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 next? Associated pages
All my
GCSE Level Chemistry Revision
notes
email doc
brown - comment? query?
See also
Salt -
sodium chloride - extraction - uses of halogens
Preparation of
hydrogen chloride & chlorine gases is described on gas preparation page
Method Ex 4
Advanced Level pre-university 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
PLEASE NOTE that these LINKS are for Advanced Level (UK 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
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