Task 1.6b Explain these spectra.
1.6a The physical properties of the elements
Elements of both groups exist as solids at room temperature due to
the strength of metallic bonding. The delocalised electrons in the structures
give each element a silvery sheen and causes them to be good electrical
conductors. They differ from other metals of the periodic table in a number of
ways: They are soft and group 1 metals can be cut with a knife. They have
low melting points, boiling points, low standard enthalpies of fusion and
vaporisation, and low densities (Li, Na and K are less dense than water).
On descending both groups, the atomic radii become bigger but the number
of delocalised electrons remain the same. The metallic bonds then become
weaker as delocalised electrons become more thinly spread as metallic radius
increases. This is why the hardness, melting and boiling points and
standard enthalpies of fusion and vaporisation decrease from top to bottom.
Densities increase down a group due to the mass of the nuclei increasing
faster than the atomic radii and therefore the atomic volumes. Group I
metals are softer and have lower melting and boiling points and standard
enthalpies of fusion and vaporisation as fewer electrons are available for
metallic bonding. Group I densities are lower than group II as the smaller
atoms in group II form a more efficiently packed structure. Open body centred cubic
structures exist for group I compared to face centred cubic and hexagonal
close packed for most of group II.
Task 1.6a
1.6b The flame colours produced by cations and electronic
transitions
Coloured flames are observed when these elements are heated. This is
because their outer electrons can be excited to higher energy levels, when
they absorb energy from the source of heat. The falling back of the electrons
to the ground state causes the emission of light which for groups I and II
happens to be in the visible region. The energy emitted is only of certain
allowed quantities which correspond to certain electronic transitions.
Lithium-
scarlet(red), Calcium- brick red, Strontium- crimson(red),
Potassium- lilac,
Barium- apple green, Sodium
- yellow
Spectral analysis: The colours are seen as a bands of coloured lines
through a spectrometer. Each element has a specific set of lines and so
can be identified in the emission spectrum even if other elements are present.
Light coming from a gas in the laboratory or from a distant star can be
analysed in this way
Task 1.6b Explain these spectra.
1.6c Group trends in ionisation energies
In both groups the trend in ionisation energy is to decrease from top to bottom.
Down a group outer electrons become further from the nucleus as they are
increasingly distant shells so are pulled
less strongly causing ionisation energy to decrease. Shielding increases
down a group which has a similar effect on ionisation energy. The
second ionisation energy of group 1 elements is much high than the first.
The second electron must be removed from an already positive ion and from
a stable full energy level closer to the nucleus. The third ionisation
energy of group 2 elements is much higher than the second for similar reasons.
Task 1.6c
1.6d Reactions of group 1 with oxygen, chlorine and water
Reactions with Oxygen
All the s block elements, readily react with air to form oxides.
All Group I metals react with oxygen to form normal oxides
(O2-).
All except Li react to form a peroxide(O22-).
Only K, Rb and Cs react to form superoxides (O2-).
4Li(s) + O2(g) ----> 2Li2O(s)
[normal oxide O2-]
2Na(s) + O2(g) -----> Na2O2(s)
[peroxide O22-]
K(s) + O2(g) -----> KO2(s)
[superoxide
O2- ]
Other peroxides and superoxides are not formed because the highly polarising
ions like Li+ cause these other oxide ions to decompose.
Reaction with Chlorine
The group I metals reduce Chlorine on heating to form a chloride.
2Na(s) + Cl2(g) ------> 2NaCl(s)
Reaction with water
Groups I metals react with cold water,
to give hydrogen and the metal hydroxide. This dissolves to give an alkaline
solution. Group I metals are kept under
oil to prevent them from reacting with water vapour in the air. The reactivity with water
increases down Group I. Therefore, Rubidium and Caesium
react very violently with water compared to Lithium, which reacts slowly;
2Na(s) + 2H2O(l) -----> 2NaOH(aq)
+ H2(g)
Task 1.6d:
Write balanced equations for the reactions of
oxygen with sodium, lithium and caesium;
chlorine with rubidium, lithium and potassium
water with lithium potassium and caesium
1.6e Reactions of group 2 with oxygen, chlorine and water
Reactions with Oxygen
All Group II metals react with oxygen to form normal oxides
(O2-) .
All except Be, Mg and Ca, react to form peroxides(O22-).
Group II metals do not form superoxides.
Peroxide and superoxides are not formed because the small high charge ions
involved have a high polarising power and so distort and break up peroxide or
superoxide ions.
e.g. 2Mg(s) + O2(g) ------> 2MgO(s)
[normal oxide O2-]
Reaction with Chlorine
The group II metals reduce Chlorine on heating to form a chloride.
Mg(s) + Cl2(g) ------> MgCl2(s)
Reaction with water
Except for Be and Mg, Groups II metals react with cold water,
Magnesium
requires steam to react vigorously as it will react very slowly with cold
water. It will form its oxide and hydrogen.
The reactivity with water
increases down Group II. Strontium and Barium are more reactive than Calcium.
Mg(s) + H2O(l) -----> H2(g)
+ MgO(s)
Task 1.6e:
Write balanced equations for the reactions of
oxygen with beryllium, calcium and barium;
chlorine with calcium, barium and strontium;
water with calcium barium and strontium.
1.6f Group 1 and 2 oxides with water and acids
The oxides are all strong bases. The oxide ion behaves as a base by
pulling a proton from water.
O2- + H2O ---> 2OH-
O22- +2H2O ---> 2OH- +
H2O2
O2- + 2H2O ---> 2OH- +
H2O2 + O2
Group I oxides are all soluble and form hydroxides. A similar reaction
is seen for group II but low solubility at the top of the group means MgO
forms very little hydroxide. BeO does not react as due to its covalent
character O2- is not fully formed.
All of the oxides react well with acids although sulphuric acid does not always
react completely because it forms insoluble sulphates with Ca, Sr and Ba.
O2- + 2H+ ---> H2O
O22- + 2H+ --->
H2O2
2O2- + 2H+ --->
H2O2
+ O2
Task 1.6f Write the ionic equations above as full balanced
equations showing specific oxides, acids or alkalis.
1.6g Oxidation numbers of group 1 and 2 elements in their
compounds
Group I metals always have an oxidation number of +1 in their compounds.
Group II metals always have an oxidation number of +2 in their compounds.
1.6h Solubility trends for group 2 sulphates and hydroxides
The solubility of the sulphates decreases down the group. Magnesium
sulphate is very soluble, barium sulphate is insoluble and is part of the test
for sulphates.
The solubility of the hydroxides increases down the group. Calcium
hydroxide is only slightly soluble in limewater but barium hydroxide is a very
soluble alkali which can be used in titrations.
1.6i The thermal stability of carbonates and nitrates
As polarising power of the cation increases, compounds become more covalent in character
and less stable to heat. Group I metal ions have low charges and large
sizes so their polarising power is low, so group I nitrates and carbonates are
more stable to heat than the corresponding Group II compounds. Carbonates and nitrates of both groups
become more thermally stable as we go down the group as cationic size increases,
and polarising power decreases. A decrease in polarising power of the cation
means the distorting effect on the anion's structure is reduced, making it more thermally
stable.
Except lithium nitrate, all the Group I metal nitrates decompose on strong heating forming nitrites and oxygen.
e.g. 2NaNO3(s) ------> 2NaNO2(s) +O2(g)
Lithium and Group II nitrates decompose to form nitrogen dioxide, oxygen and the thermally stable oxides.
e.g. 2Ca(NO3)2(s) -----> 2CaO(s) + 4NO2(g) + O2(g)
4LiNO3(s) -----> 2Li2O(s) + 4NO2(g) + O2(g)
The carbonates of all s block metals in group II and lithium all decompose to form an oxide and carbon dioxide.
e.g. Li2CO3(s) ----> Li2O(s) + CO2(g)
MgCO3(s) ----> MgO(s) + CO2(g)
Carbonates of Na, K and Rb do not decompose on heating at normal temperatures.