Group 2 Revision Guide Flashcards
Melting points in group 2
Meiting points decrease down the group. The metallic bonding weakens as the atomic size increases. The distance between the positive ions and delocalized electrons increases. Therefore the electrostatic attractive forces between the positive ions and the delocalized electrons weaken.
1st Ionisation Energy
The outermost electrons are held more weakly because they are successively further from the nucleus in additional shells.
In addition, the outer shell electrons become more shielded from the attraction of the nucleus by the repulsive force of inner shell electrons.
Reactivity
Reactivity of group 2 metals increases down the group
Mg Reactions with oxygen
The group 2 metals will burn in oxygen.
Mg burns with a bright white flame. The
MgO appears as a white powder.
2 Mg + O2 -> 2 MgO
MgO is a white solid with a high melting
point due to its ionic bonding.
How is MgO formed
Magnesium will also react slowly in oxygen without a flame. Magnesium ribbon will often have a thin layer of magnesium oxide on it formed by the reaction with oxygen in the air.
2 Mg + O2 —> 2 MgO
What is the effect of magnesium oxide and how do we remove this?
The magnesium oxide needs to be removed by emery paper before doing reactions with magnesium ribbon. If testing for reaction rates with Mg and acid, an un-cleaned Mg ribbon would give a false result because both the Mg and MgO would react but at different rates.
Mg + 2 HCl —> MgCl2 + H2
MgO + 2 HCl —> MgCl2 + H2O
Magnesium reactions with steam
Magnesium reacts in steam to produce magnesium oxide and hydrogen the magnesium would be with a bright white flame. The MgO appears as a white powder.
Mg (s) + H2O(g) —> MgO(s) + H2(g)
Magnesium reacting with warm water
Magnesium will also react with warm water giving a different magnesium hydroxide product
This is a much slower reaction done the reaction with steam as there is no flame
Mg + 2H2O —> Mg(OH)2 + H2
Group 2 metals reacting with cold water and their equasions
The other group 2 metals will react with cold water with increasing vigour down the group to form hydroxides.
Ca + 2 H2O (l) —> Ca(OH)2 (aq) + H2 (g)
Sr + 2 H2O (l) —> Sr(OH)2 (aq) + H2 (g)
Ba + 2 H2O (l) —> Ba(OH)2 (aq) + H2 (g)
Observations with water reactions
fizzing, (more vigorous down group)
the metal dissolving, (faster down group)
the solution heating up (more down group)
with calcium a white precipitate appearing (less precipitate forms down group with other metals)
Why is titanium useful
Titanium is a very useful metal because it is abundant, has a low density and is corrosion resistant – it is used for making strong, light alloys for use in aircraft.
What is titanium extracted with
Titanium cannot be extracted with carbon because titanium carbide (TiC) it is formed rather than titanium.
Titanium cannot be extracted by electrolysis because it has to be very pure.
Titanium is extracted by reaction with a more reactive metal (e.g. Magnesium).
Equasions for titanium extraction
TiO2 + 2 Cl2 + 2 C —> TiCl4 + 2 CO
TiCl4 + 2Mg —> Ti + 2 MgCl2
Steps in extracting titanium
- TiO2 (solid) is converted to TiCl4 (liquid) at 900 degrees:
- The TiCl4 is purified by fractional distillation in an argon atmosphere.
- The Ti is extracted by Mg in an argon atmosphere at 500 degrees
Why is titanium expensive
Titanium is expensive because:
The expensive cost of the magnesium
This is a batch process which makes it expensive because the process is slower (having to fill up and empty reactors takes time) and requires more labour and the energy is lost when the reactor is cooled down after stopping
- The process is also expensive due to the argon, and the need to
remove moisture (because TiCl4 is susceptible to hydrolysis). - High temperatures required in both steps
Why is TiO2 converted to TiCl4
TiO2 is converted to TiCl4 as it can be purified by fractional distillation, TiCl4 being molecular (liquid at room temperature) rather than ionic like TiO2 (solid at room temperature).
What is calcium oxide used for
Calcium oxide can be used to remove SO2 from the waste gases from furnaces (e.g. coal fired power stations) by flue gas desulfurisation. The gases pass through a scrubber containing basic calcium oxide which reacts with the acidic sulfur dioxide in a neutralisation reaction.
SO2 + CaO —> CaSO3
Trend of solubility of hydroxides
Group II hydroxides become more soluble down the group.
All Group II hydroxides when not soluble appear as white precipitates.
Insolubility of mgoh2
Magnesium hydroxide is classed as insoluble in water
Simplest ionic equation for the formation of Mg(OH)2(s)
Mg2+(aq) + 2OH-(aq) —> Mg(OH)2(s).
acidic soils.
A suspension of magnesium hydroxide in water will appear slightly alkaline (pH 9) so some hydroxide ions must therefore have been produced by a very slight dissolving.
Uses of magnesium hydroxide in mediicne
Magnesium hydroxide is used in medicine (in suspension as milk of magnesia) to neutralise excess acid in the stomach and to treat constipation.
Mg(OH)2 + 2HCl —> MgCl2 + 2H2O
It is safe to use because it is so weakly alkaline. It is preferable to using calcium carbonate as it will not produce carbon dioxide gas.
Solubility of calcium hydroxide
A suspension of calcium hydroxide in water will appear more alkaline (pH 11) than magnesium hydroxide as it is more soluble so there will be more hydroxide ions present in solution.
An aqueous solution of calcium hydroxide is called lime water and can be used a test for carbon dioxide. The limewater turns cloudy as white calcium carbonate is produced.
Ca(OH)2 (aq) + CO2 (g) —> CaCO3 (s) + H2O(l)
What is calcium hydroxide used for
Calcium hydroxide is classed as partially
soluble in water and will appear as a white
precipitate It is used in agriculture to neutralise acidic soils.
Solubility of barium hydroxide
Barium hydroxide would easily dissolve in
water. The hydroxide ions present would make the solution strongly alkaline.
Ba(OH)2 (s) + aq —> Ba2+ (aq) + 2OH-(aq)
Trend of solubility of sulfates
Group II sulfates become less soluble down the group. BaSO4 is the least soluble
An ionic and full equation for the formation of the precipitate
An equation for the formation of the precipitate can be written as a full equation or simplest ionic equation.
Full equation : SrCl2(aq) + Na2SO4 (aq) —> 2NaCl (aq) + SrSO4 (s)
Ionic equation: Sr2+ (aq) + SO42-(aq) —> SrSO4 (s).
How is barium sulfate used in medicine
BaSO4 is used in medicine as a ‘Barium meal’ given to patients who need x-rays of their intestines. The barium absorbs the x-rays and so the gut shows up on the x-ray image.
Even though barium compounds are toxic, it is safe to use here because barium sulfate’s low solubility means it is not absorbed into the blood.
Reactions of barium sulfate
If barium metal is reacted with sulfuric acid it will only react slowly, as the insoluble barium sulfate produced will cover the surface of the metal and act as a barrier to further attack.
Ba + H2SO4 —> BaSO4 + H2
The same effect happens to a lesser extent with metals going up the group as the solubility of the sulfates increases. The same effect does not happen with other acids like hydrochloric or nitric as they form soluble group 2 salts.
What is used to test for sulfate ion
BaCl2 solution acidified with hydrochloric acid is used as a reagent to test for sulfate ions.
If acidified barium chloride is added to a solution that contains sulfate ions a white precipitate of barium sulfate forms.
Ionic equasion
Ba2+ (aq) + SO42-(aq) —> BaSO4 (s)
Why is hcl needed to test for sulfate ions aswell
The hydrochloric acid is needed to react with carbonate impurities that are often found in salts which would form a white barium carbonate precipitate and so give a false result.
You could not use sulfuric acid because it contains sulfate ions and so would give a false positive result.
2HCl + Na2CO3 —> 2NaCl + H2O + CO
Fizzing due to CO2 would be observed if a carbonate was present
How can insoluble salts be made
Insoluble salts can be made by mixing appropriate solutions of ions so that a precipitate is formed
barium nitrate (aq) + sodium sulfate (aq) —> barium sulfate (s) + sodium nitrate (aq)
These are called precipitation reactions. A precipitate is a solid.
Examples of soluble salts
All sodium, potassium and ammonium salts
All nitrates
Most chlorides, bromides, iodides
Most sulfates Sodium, potassium and ammonium carbonates
Sodium, potassium and ammonium hydroxides
Insoluble salt examples
Silver, lead chlorides, bromides iodides
Lead, strontium and barium sulfates Most other carbonates
Most other hydroxides
How is salt removed from insoluble salts
When making an insoluble salt, normally the salt would be removed by filtration, washed with distilled water to remove soluble impurities and then dried on filter paper.
What are spectator ions
Ions that are
Not changing state
Not changing oxidation number
