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Chemistry Unit 1.6: Groups 1 and 2
The atomic radius increases as you go down group 1 and group 2. This is because the radius of an
atom is governed by two factors:
the number of layers of electrons around the nucleus
the pull the outer electrons feel from the nucleus.
As all of group 1 feel a net pull of +1 and group 2 a net pull of +2 (nuclear charge minus no. inner
electrons), the only factor which is going to affect the size of the atom is therefore the number of
layers of inner electrons which have to be fitted in around the atom. Obviously, the more layers of
electrons you have, the more space they will take up - electrons repel each other. That means that
the atoms are bound to get bigger as you go down both groups.
Also the melting points decrease as you go down both groups and the densities increase as you go
Sodium - orange/yellow
Potassium - lilac
Rubidium Reddish violet
Caesium - Blue/violet
Strontium - Red
Barium Apple green
Flame colours originate from when the heat energy in the flame promotes and electron in a lower
energy level to a higher shell, the atom is now said to be in an excited state. When the electron
moves back down, the energy is released as a particular wavelength of light, which determines the
First Ionisation Energies
The first ionisation energy is the energy needed to remove the most loosely held electron from each
of one mole of gaseous atoms to make one mole of singly charged gaseous ions. The first ionisation
energies of groups 1 and 2 both fall as you go down the group this is because ionisation energy is
the charge on the nucleus,
the amount of screening by the inner electrons,
the distance between the outer electrons and the nucleus.
As you go down the Groups, the increase in nuclear charge is exactly offset by the increase in the
number of inner electrons, just like atomic radius. They all feel the same net pull in group 1 of +1 and
in group 2 of +2.
However, as you go down the Group, the distance between the nucleus and the outer electrons
increases and so they become easier to remove - the ionisation energy falls.
Group 1 reactions with O2, Cl2 and H2O
Lithium - Lithium burns with a strongly red-tinged flame if heated in air. It reacts with oxygen in
the air to give white lithium oxide. With pure oxygen, the flame would simply be more intense.
Sodium - Small pieces of sodium burn in air with often little more than an orange glow. Using
larger amounts of sodium or burning it in oxygen gives a strong orange flame. You get a white
solid mixture of sodium oxide and sodium peroxide.
Potassium - Small pieces of potassium heated in air tend to just melt and turn instantly into a
mixture of potassium peroxide and potassium superoxide without any flame being seen. Larger
pieces of potassium burn with a lilac flame.
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Rubidium and caesium - Both metals catch fire in air and produce superoxides, RbO2 and CsO2.
The equations are the same as the equivalent potassium one.
All the group 1 metals react with chlorine in the same way they do with the oxides except they only
ever form the simple chloride with the formula XCl.
All react vigorously to form ionic chlorides formula MCL and dissolve in water to produce hydrated
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Reaction with dilute acids - These simple oxides all react with an acid to give a salt and water.
For example, sodium oxide will react with dilute hydrochloric acid to give colourless sodium
chloride solution and water.
Reaction with water - If the reaction is done ice cold (and the temperature controlled so that it
doesn't rise even though these reactions are strongly exothermic), a solution of the metal
hydroxide and hydrogen peroxide is formed.…read more
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If "X" represents any one of the elements:
The carbonates become more stable to heat as you go down the Group.
All the nitrates in this Group undergo thermal decomposition to give the metal oxide, nitrogen
dioxide and oxygen.
Again, if "X" represents any one of the elements:
The nitrates also become more stable to heat as you go down the Group.
A small 2+/1+ ion has a lot of charge packed into a small volume of space.…read more