Download Groups 2 and 7

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Physical properties
The arrangement of the periodic table is
such that trends can be analysed both
across a period and down a group.
Group 2 of the periodic table is shown
here. Trends that can be analysed
down the group include atomic
radius, first ionization energy and
melting point.
Elements in the same group also
undergo similar chemical reactions.
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Trend in atomic radius
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Explaining the trend in atomic radius
The atomic radius of the
elements increases down
Element
group 2 from beryllium to barium.
The number of protons
increases down the group;
however, so does the number
of shielding electrons. Effective
nuclear charge therefore
remains approximately
constant.
Atomic
radius
(nm)
beryllium
0.112
magnesium
0.145
calcium
0.194
strontium
0.219
barium
0.253
The increase in radius is due to higher principle energy
levels being filled, whose orbitals are located further from
the nucleus.
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Trend in first ionization energy
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First ionization energies in group 2
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Trend in melting points
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Explaining the trend in melting points
The melting points of the
elements decrease down
group 2, with the exception
of magnesium to calcium.
A metal’s melting point depends
on the strength of its metallic
bonds. This decreases down the
group because the atomic radius
increases, resulting in a weaker
attraction between the nucleus
and delocalized electrons.
Element
Melting
point (K)
beryllium
1560
magnesium
923
calcium
1115
strontium
1050
barium
1000
The melting point of magnesium is lower than expected due
to variation in how its atoms pack in the metallic crystal.
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Physical properties summary
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First ionization energy of group 2 metals
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Solubilities of group 2 hydroxides
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Solubilities of group 2 hydroxides
The solubility of the
group 2 hydroxides
increases down the
group. Magnesium
hydroxide is considered
to be sparingly soluble
and the hydroxides of
the lower members of
the groups are all
considered to be soluble.
Group 2
hydroxide
Solubility
Mg(OH)2
sparingly soluble
Ca(OH)2
slightly soluble
Sr(OH)2
soluble
Ba(OH)2
soluble
As the solubility of the group 2 hydroxides increases, so
does the pH of the solutions formed. This is because the
more of the hydroxide that dissolves, the greater the
concentration of hydroxide ions (OH-) in the solution formed.
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Applications of group 2 hydroxides
A suspension of magnesium hydroxide is commonly called
milk of magnesia. It is used in medicine as a laxative and to
relieve acid indigestion.
Calcium hydroxide, also
called slaked lime, is
used in agriculture to raise
the pH of soils. Soil pH is
an important factor in
agriculture.
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Solubilities of group 2 sulfates
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Solubilities of group 2 sulfates
The solubility of the group 2 sulfates decreases down the
group. Magnesium and calcium sulfate are considered to
be soluble, whereas strontium and barium sulfate are
considered to be insoluble.
Group 2
hydroxide
Solubility
MgSO4
soluble
CaSO4
slightly soluble
SrSO4
insoluble
BaSO4
insoluble
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Note that this
decrease in solubility
down the group is
the opposite of the
trend for the
solubility of the
group 2 hydroxides.
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Testing for sulfate ions
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Applications of group 2 sulfates
Barium sulfate is used as a
radiocontrast agent to help
take X-ray images of the
digestive system. It is
sometimes known as a
‘barium meal’.
Barium sulfate is insoluble, so
is not absorbed by the body
when swallowed. However,
barium is a very good absorber
of X-rays and it helps to define
structures of the digestive
system to aid in diagnosis.
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Applications of group 2 compounds
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Chemical properties summary
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Reaction with oxygen
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Flame tests
When group 2 metals are burned in oxygen, coloured
flames are produced. This is due to the presence of metal
ions. Flame tests exploit this fact.
The presence of certain metal ions can be identified by
noting the characteristic flame colour that results from
burning. The colours for group 2 metal ions are:
magnesium – bright white
calcium – brick red/orange
strontium – red/crimson
barium – pale green/yellow-green
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Explaining flame tests
When heated, some electrons in an atom or ion are excited
to higher energy levels. When they fall back to their initial
levels, energy is emitted; sometimes seen as visible light.
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light
heat
energy
Electrons may be excited by
different amounts into
different energy levels and
drop back at different times.
The colour of the flame is a
combination of all these
energy emissions.
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Flame test colours
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Redox reaction with oxygen
When group 2 metals react with oxygen, they form the
metal oxide. For example:
2Mg(s) + O2(g)  2MgO(s)
0
0
+2 -2
oxidation
states
The oxidation state of magnesium has increased from 0 in
its elemental form to +2 when it is in magnesium oxide.
This means the magnesium has been oxidized.
The oxidation state of oxygen has decreased from 0 in its
elemental form to -2 when it is in magnesium oxide. This
means the oxygen has been reduced.
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Redox reaction with chlorine
When group 2 metals react with chlorine, they form the
metal chloride. For example:
Ca(s) + Cl2(g) CaCl2(s)
0
0
+2 -1
oxidation
states
The oxidation state of calcium has increased from 0 in its
elemental form to +2 when it is in calcium chloride. This
means the calcium has been oxidized.
The oxidation state of chlorine has decreased from 0 in its
elemental form to -1 when it is in calcium chloride. This
means the chlorine has been reduced.
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Reaction with water
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Redox reaction with water
When group 2 metals react with water they form the metal
hydroxide and hydrogen gas. For example:
Sr(s) + 2H2O(l) → Sr(OH)2(aq) + H2(g)
0
+1
+2
0
oxidation
states
The oxidation state of strontium has increased from 0 in
its elemental form to +2 when it is in strontium hydroxide.
This means the strontium has been oxidized.
The oxidation state of hydrogen has decreased from +1 in
water to 0 when it is in its elemental form. The means the
hydrogen has been reduced.
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Explaining the trend in reactivity
The reactivity of the elements
down group 2 from beryllium to
barium increases.
Mg
This is because it is
successively easier to remove
electrons to form the 2+ ion.
Ca
Although increased shielding
cancels the increased nuclear
charge down the group, the
increase in atomic radius results
in a decrease in the attractive
force between the outer
electrons and the nucleus.
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Sr
Ba
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Reaction of oxides with water
When group 2 metal oxides react with water they form the
metal hydroxide. For example:
SrO(s) + H2O(l)  Sr(OH)2(aq)
Similar to the reaction between the metal and water, the
resulting solution has high pH due to the hydroxide ions
from the metal hydroxide. Reactivity is as follows:
Oxide
beryllium
magnesium
calcium
Reaction
does not react
reacts slowly to form
alkaline suspension
reacts to form alkaline suspension
strontium, barium react to form alkaline solutions
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Decomposition of group 2 carbonates
When heated, the group 2 metal carbonates decompose
to form the metal oxide and carbon dioxide gas. Splitting
compounds using heat is called thermal decomposition.
MCO3(s)  MO(s) + CO2(g)
The group 2 carbonates become more stable to thermal
decomposition going down the group:
magnesium carbonate: MgCO3
increasing
stability
calcium carbonate: CaCO3
strontium carbonate: SrCO3
barium carbonate: BaCO3
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Decomposition of group 2 nitrates
Thermal decomposition of group 2 metal nitrates forms
the metal oxide, nitrogen dioxide and oxygen.
2M(NO3)2(s)  2MO(s) + 4NO2(g) + O2(g)
Like the group 2 metal carbonates, the nitrates become
more stable to thermal decomposition down the group.
magnesium nitrate: Mg(NO3)2
increasing
stability
calcium nitrate: Ca(NO3)2
strontium nitrate: Sr(NO3)2
barium nitrate: Ba(NO3)2
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Explaining the trend in thermal stability
Metal ions become larger down group 2 but have the same
charge. This means their charge density is reduced.
A metal ion with a high
charge density has
strong polarizing power.
It can therefore polarize
the carbonate ion,
making it more likely to
split into O2- and CO2
when heated.
polarization
A metal ion with a low charge density has weak polarizing
power, meaning the carbonate ion is less polarized and
therefore more thermally stable.
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Equations for reactions
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Stability of group 2 carbonates
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Glossary
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What’s the keyword?
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Multiple-choice quiz
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What are the halogens?
The halogens are the elements in Group 7 of the
periodic table.
The name halogen comes from the Greek
words for salt-making.
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Physical properties of halogens
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Trends in boiling point
Halogen molecules increase in size down the group. This
leads to greater van der Waals forces between molecules,
increasing the energy needed to separate the molecules
and therefore higher melting and boiling points.
van der
Waals forces
fluorine
atomic radius = 42 × 10-12 m
boiling point = -118 °C
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iodine
atomic radius = 115 × 10-12 m
boiling point = 184 °C
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Trends in electronegativity
Electronegativity of the halogens decreases down the group
due to an increase in atomic radius.
Increased nuclear charge has no significant effect because
there are more electron shells and more shielding. Iodine
atoms therefore attract electron density in a covalent bond
less strongly than fluorine.
fluorine
atomic radius = 42 × 10-12 m
electronegativity = 4.0
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iodine
atomic radius = 115 × 10-12 m
electronegativity = 2.5
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Astatine
The name astatine comes from the Greek word for unstable.
Astatine exists in nature in only very tiny
amounts. It is estimated that only 30 grams of
astatine exist on Earth at any one time. This is
because it is radioactive, and its most stable
isotope (210At) has a half-life of only 8 hours.
It was first made artificially in 1940, by bombarding 209Bi with
a-radiation. What do you predict for these properties of
astatine?

colour

state at room temperature

electronegativity.
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Halogens: true or false?
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Reactions of the halogens
Halogens react with metals such as sodium and iron:
halogen + sodium  sodium halide
They also react with non-metals such as hydrogen:
halogen + hydrogen  hydrogen halide
They also take part in displacement reactions with halide
ions, such as the reaction that is used to make bromine from
potassium bromide in seawater:
potassium
potassium
chlorine +
 bromine +
bromide
chloride
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Reaction with iron
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Reactions with hydrogen
The halogens react with hydrogen gas to product hydrogen
halides. For example:
Cl2(g) + H2(g)  2HCl(g)

Chlorine and hydrogen explode in bright sunlight but
react slowly in the dark.

Bromine and hydrogen react slowly on heating with
a platinum catalyst.

Iodine combines partially and very slowly with
hydrogen, even on heating.
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Redox reactions of halogens
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What is the reactivity of the halogens?
The reactions of the halogens with iron and hydrogen show
that their reactivity decreases down the group.
Halogen Reaction with iron wool
chlorine Iron wool burns and
glows brightly.
Reaction with
hydrogen
Explodes in
sunlight, reacts
slowly in the dark.
bromine Iron wool glows but less
Reacts slowly on
brightly than with chlorine. heating with catalyst.
iodine
Iron wool has a very
slight glow.
Reacts partially
and very slowly.
How do you think fluorine and astatine would react with
iron wool and hydrogen?
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Electron structure and reactivity
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Halogen displacement reactions
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Halogen displacement reactions
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Halogen displacement reactions
Halogen displacement reactions are redox reactions.
Cl2 + 2KBr  2KCl + Br2
To look at the transfer of electrons in this reaction, the
following two half equations can be written:
Cl2 + 2e-  2Cl-
2Br-  Br2 + 2e-
What has been oxidized and what has been reduced?

Chlorine has gained electrons, so it is reduced to Cl- ions.

Bromide ions have lost electrons, so they have been
oxidized to bromine.
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In displacement reactions between
halogens and halides, the halogen
acts as an oxidizing agent.
This means that the halogen:

oxidizes the halide ion to the
halogen

gains electrons

is reduced to form the halide ion.
What is the order of oxidizing
ability of the halogens?
increasing oxidizing ability
Oxidizing ability of halogens
fluorine
chlorine
bromine
iodine
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Oxidizing ability of halogens
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Chlorine and disproportionation
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Reaction of chlorine with water
Chlorine is used to purify water supplies
because it is toxic to bacteria, some of
which can cause disease. Adding it to
water supplies is therefore beneficial for
the population.
However, chlorine is also toxic to humans,
so there are risks associated with gas leaks
during the chlorination process. There is
also a risk of the formation of chlorinated
hydrocarbons, which are also toxic.
Chlorination of drinking water raises questions about individual
freedom because it makes it difficult for individuals to opt out.
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Bleach and the chlorate(I) ion
Household bleach commonly contains the chlorate(I) ion,
ClO-, in the form of sodium chlorate(I), NaOCl.
The chlorate(I) ion behaves as an oxidizing agent. It oxidizes
the organic compounds in food stains, bacteria and dyes.
ClO- + H2O + 2e-  Cl- + 2OHHow many electrons are needed to balance this equation?
Has the chlorine been oxidized or reduced in the reaction?
The chlorine has been reduced because it has gained
electrons. Its oxidation state has decreased from +1 in
ClO- to –1 in Cl-.
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Redox reactions of chlorate ions
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Halides
When halogens react with metals, they form compounds
called halides. Many naturally-occurring halides have
industrial, household and medical applications.
Halide
Formula Uses
caesium chloride
CsCl
Extraction and
separation of DNA
sodium
hexafluoroaluminate
NaAlF6
Electrolysis of
aluminium oxide
titanium(IV) chloride
TiCl4
Extraction of titanium
lithium iodide
LiI
Electrolyte in batteries
potassium bromide
KBr
Epilepsy treatment in
animals
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Identifying halide ions
Halides can be identified by their reaction with acidified
silver nitrate solution to form silver halide precipitates.
potassium
chloride
+
silver
nitrate

potassium
nitrate
+
silver
chloride
KCl(aq) + AgNO3(aq)  KNO3(aq)
+ AgCl(s)
Silver chloride has a low solubility
in water, so it forms a white
precipitate: the positive result in
the test for chloride ions.
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Identifying halide ions
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Identifying halide ions
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Uses of halides in photography
Silver halides are used in photography.
Photographic film coated with a silver halide is exposed to
light, causing the halide to decompose to form silver. This
appears as a black precipitate on the photographic film.
Ag+ + e-  Ag
light
mask
paper
coated in
silver halide
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silver
precipitate
white paper
under
mask
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William Fox Talbot
William Fox Talbot (1800–1877) was a British scientist and
mathematician. He was one of the key figures in the
development of the use of silver halides in photography.
A French scientist called Louis Daguerre developed the
use of silver halides on copper plates. These were effective
at producing prints, but could only be used once.
Fox Talbot adapted the process by
removing any unreacted silver
halide by washing with sodium
thiosulfate solution. This meant that
the print could be used repeatedly
in the way that photographic
negatives can be today.
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Hydrogen halides
The hydrogen
halides are
colourless gases
at room
temperature.
Hydrogen halide
Boiling point (°C)
HF
20
HCl
-85
HBr
-67
HI
-35
Hydrogen fluoride has an
unexpectedly high boiling point
compared to the other
hydrogen halides. This is due to
hydrogen bonding between the
H–F molecules.
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A substance that donates electrons
in a reaction (i.e. is oxidized) is a
reducing agent because it
reduces the other reactant.
The larger the halide ion, the easier
it is for it to donate electrons and
therefore the more reactive it is.
This is because its outermost
electrons are further from the
attraction of the nucleus and more
shielded from it by other electrons.
The attraction for the outermost
electrons is therefore weaker.
increasing reducing ability
Halides as reducing agents
fluoride
chloride
bromide
iodide
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Halides: true or false?
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Sodium halides and sulfuric acid
The sodium halides react with concentrated sulfuric acid.
During this reaction two things can
happen to the sulfuric acid. It can

be reduced

act as an acid.
The reactions of sodium halides with concentrated sulfuric
acid demonstrate the relative strengths of the halide ions
as reducing agents.
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Sodium halides and sulfuric acid
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Sodium halides and sulfuric acid
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Oxidation states
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Sodium halides and sulfuric acid
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Glossary
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What’s the keyword?
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Multiple-choice quiz
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