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Transcript of 1 of 41© Boardworks Ltd 2009. 2 of 41© Boardworks Ltd 2009.
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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 group 2 from beryllium to barium.
The increase in radius is due to higher principle energy levels being filled, whose orbitals are located further from the nucleus.
ElementAtomicradius(nm)
beryllium
magnesium
calcium
strontium
barium
0.112
0.145
0.194
0.219
0.253
The number of protons increases down the group; however, so does the number of shielding electrons. Effective nuclear charge therefore remains approximately constant.
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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. beryllium
magnesium
calcium
strontium
barium
Element Melting point (K)
1560
923
1115
1050
1000
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.
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.
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.
Mg(OH)2
Ca(OH)2
Sr(OH)2
Ba(OH)2
Group 2hydroxide
Solubility
sparingly soluble
slightly soluble
soluble
soluble
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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.
MgSO4
CaSO4
SrSO4
BaSO4
Group 2hydroxide
Solubility
soluble
slightly soluble
insoluble
insoluble
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.
magnesium – bright white
calcium – brick red/orange
strontium – red/crimson
barium – pale green/yellow-green
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:
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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.
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.
heatlight
energ
y
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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)
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.
0 +20 -2oxidation
states
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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 -1oxidation
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)
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.
0 +2+1 0oxidation
states
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Explaining the trend in reactivity
The reactivity of the elements down group 2 from beryllium to barium increases.
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.
This is because it is successively easier to remove electrons to form the 2+ ion.
Mg
Ca
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:
beryllium
magnesium
calcium
strontium, barium
does not reactreacts slowly to form alkaline suspension
reacts to form alkaline suspension
react to form alkaline solutions
Oxide Reaction
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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.
magnesium carbonate: MgCO3
calcium carbonate: CaCO3
strontium carbonate: SrCO3
barium carbonate: BaCO3
increasing stability
The group 2 carbonates become more stable to thermal decomposition going down the group:
MCO3(s) MO(s) + CO2(g)
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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
calcium nitrate: Ca(NO3)2
strontium nitrate: Sr(NO3)2
barium nitrate: Ba(NO3)2
increasing stability
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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.
fluorineatomic radius = 42 × 10-12
mboiling point = -118 °C
iodine atomic radius = 115 × 10-12
mboiling point = 184 °C
van der Waals forces
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Trends in electronegativity
Electronegativity of the halogens decreases down the group due to an increase in atomic radius.
fluorineatomic radius = 42 × 10-12
melectronegativity = 4.0
iodineatomic radius = 115 × 10-12
melectronegativity = 2.5
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.
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Astatine
The name astatine comes from the Greek word for unstable.
It was first made artificially in 1940, by bombarding 209Bi with -radiation. What do you predict for these properties of astatine?
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.
electronegativity.
state at room temperature
colour
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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:
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:
halogen + hydrogen hydrogen halide
They also react with non-metals such as hydrogen:
halogen + sodium sodium halide
chlorine +potassiumbromide
potassiumchloride
bromine +
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Reaction with iron
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Reactions with hydrogen
Chlorine and hydrogen explode in bright sunlight but react slowly in the dark.
The halogens react with hydrogen gas to product hydrogen halides. For example:
Cl2(g) + H2(g) 2HCl(g)
Iodine combines partially and very slowly with hydrogen, even on heating.
Bromine and hydrogen react slowly on heating with a platinum catalyst.
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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.
How do you think fluorine and astatine would react with iron wool and hydrogen?
Iron wool burns and glows brightly.
Iron wool has a very slight glow.
Iron wool glows but less brightly than with chlorine.
chlorine
bromine
iodine
Halogen Reaction with iron wool
Reaction with hydrogen
Explodes in sunlight, reacts slowly in the dark.
Reacts slowly on heating with catalyst.
Reacts partially and very slowly.
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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:
Chlorine has gained electrons, so it is reduced to Cl- ions.
What has been oxidized and what has been reduced?
2Br- Br2 + 2e-Cl2 + 2e- 2Cl-
Bromide ions have lost electrons, so they have been oxidized to bromine.
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Oxidizing ability of halogens
fluorine
incr
easi
ng
oxi
diz
ing
ab
ility
iodine
bromine
chlorine
In displacement reactions between halogens and halides, the halogen acts as an oxidizing agent.
This means that the halogen:
What is the order of oxidizing ability of the halogens?
is reduced to form the halide ion.
gains electrons
oxidizes the halide ion to the halogen
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Oxidizing ability of halogens
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Chlorine and disproportionation
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Reaction of chlorine with water
Chlorination of drinking water raises questions about individual freedom because it makes it difficult for individuals to opt out.
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.
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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.
ClO- + H2O + Cl- + 2OH-
The chlorine has been reduced because it has gained electrons. Its oxidation state has decreased from +1 in ClO- to –1 in Cl-.
How many electrons are needed to balance this equation?
The chlorate(I) ion behaves as an oxidizing agent. It oxidizes the organic compounds in food stains, bacteria and dyes.
Has the chlorine been oxidized or reduced in the reaction?
2e-
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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.
caesium chloride
sodiumhexafluoroaluminate
titanium(IV) chloride
lithium iodide
potassium bromide
Halide Formula Uses
CsCl
NaAlF6
TiCl4LiI
KBr
Extraction and separation of DNA
Electrolysis of aluminium oxide
Extraction of titanium
Electrolyte in batteries
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.
Silver chloride has a low solubility in water, so it forms a white precipitate: the positive result in the test for chloride ions.
KCl(aq) + AgNO3(aq) KNO3(aq)
+ AgCl(s)
potassium
chloride
silver chloride
+potassium
nitratesilver
nitrate+
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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.
Ag+ + e- Ag
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.
light
mask
paper coated in
silver halide
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.
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.
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.
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Hydrogen halides
The hydrogen halides are colourless gases at room temperature.
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.
Hydrogen halide Boiling point (°C)
HF
HCl
HBr
HI
20
-85
-67
-35
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Halides as reducing agents
A substance that donates electrons in a reaction (i.e. is oxidized) is a reducing agent because it reduces the other reactant.
fluoride
incr
easi
ng
red
uci
ng
ab
ility
iodide
bromide
chloride
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.
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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.
The reactions of sodium halides with concentrated sulfuric acid demonstrate the relative strengths of the halide ions as reducing agents.
During this reaction two things can happen to the sulfuric acid. It can
act as an acid.
be reduced
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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