Group 2- The Alkaline Earth Metals

A presentation by

Madeeha Durrani

Group 2 - The Alkaline Earth Metals

The elements of Group 2, the Alkaline Earth Metals, are: •   symbol electron configuration • beryllium Be [He]2s2 • magnesium Mg [Ne]3s2 • calcium Ca [Ar]4s2 • strontium Sr [Kr]5s2 • Barium Ba [Xe]6s2 • radium Ra [Rn]7s2

The last element, radium, is radioactive and will not be considered here. Appearance • The Group 2 elements are all metals with a shiny, silvery-white colour. General Reactivity • The alkaline earth metals are high in the reactivity series of metals, but not as high as the alkali

metals of Group 1. • Occurrence and Extraction • These elements are all found in the Earth’s crust, but not in the elemental form as they are so

reactive. Instead, they are widely distributed in rock structures. The main minerals in which magnesium is found are carnellite, magnesite and dolomite. Calcium is found in chalk, limestone, gypsum and anhydrite. Magnesium is the eighth most abundant element in the Earth’s crust, and calcium is the fifth.

Physical Properties

Atomic radii

• You can see that the atomic radius increases as you go down the Group. Notice that beryllium has a particularly small atom compared with the rest of the Group.

• Explaining the increase in atomic radius

• The radius of an atom is governed by

• the number of layers of electrons around the nucleus

• the pull the outer electrons feel from the nucleus.

Compare beryllium and magnesium:

• Be 1s2 2s2

• Mg 1s2 2s2 2p6 3s2

Trends in First Ionisation Energy

• 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

                                                                                               

Notice that first ionisation energy falls as you go down the group.

First ionization energy

Explaining the decrease in first ionisation energy Ionisation energy is governed by 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 Group, the increase in nuclear charge

is exactly offset by the increase in the number of inner electrons. Just as when we were talking about atomic radius further up this page, in each of the elements in this Group, the outer electrons feel a net attraction of 2+ from the centre.

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.

Electronegativity• Electronegativity is a measure of

the tendency of an atom to attract a bonding pair of electrons. It is usually measured on the Pauling scale, on which the most electronegative element (fluorine) is given an electronegativity of 4.0

• • All of these elements have a low

electronegativity. (Remember that the most electronegative element, fluorine, has an electronegativity of 4.0.) Notice that electronegativity falls as you go down the Group. The atoms become less and less good at attracting bonding pairs of electrons.

Explaining the decrease in electronegativity• Imagine a bond between a magnesium

atom and a chlorine atom. Think of it to start with as a covalent bond - a pair of shared electrons. The electron pair will be dragged towards the chlorine end because there is a much greater net pull from the chlorine nucleus than from the magnesium one.

• The electron pair ends up so close to the

chlorine that there is essentially a transfer of an electron to the chlorine - ions are formed.

• The large pull from the chlorine nucleus is why chlorine is much more electronegative than magnesium is.

• As the metal atoms get bigger, any bonding pair gets further and further away from the metal nucleus, and so is less strongly attracted towards it. In other words, as you go down the Group, the elements become less electronegative.

Trends in melting and boiling points

Melting and Boiling points

Melting point and boiling point decreases steadily down the group. This is because the ionic radii increase down the group increasing charge separation between the metal cations of the lattice and the free delocalised electrons. This weakens the electrical attractive bonding force and so less thermal KE is needed to weaken the lattice to the 'collapse point' i.e. melting. BUT the situation is not as simple as might be expected, e.g. the metal ions do not always have the same crystal lattice packing arrangement

Chemical Properties

Reaction with water

The Group 2 metals become more reactive towards water as you go down the Group.

BerylliumBeryllium has no reaction with water or steam even at red heat.

MagnesiumMagnesium burns in steam to produce magnesium oxide andhydrogen.

Very clean magnesium has a very slight reaction with cold water. The reaction soon stops because the magnesium hydroxide formed is almost insoluble in water and forms a barrier on the magnesium preventing further reaction.

Calcium, strontium and barium These all react with cold water with increasing vigour to give the

metal hydroxide and hydrogen. Strontium and barium have reactivities similar to lithium in Group 1 of the Periodic Table.

The equation for the reactions of any of these metals would be:

The hydroxides aren't very soluble, but they get more soluble as you go down the Group. The calcium hydroxide formed shows up mainly as a white precipitate (although some does dissolve). You get less precipitate as you go down the Group because more of the hydroxide dissolves in the water.

Explaining the trend in reactivity Beryllium as a special case There is an additional reason for the lack

of reactivity of beryllium compared with the rest of the Group. Beryllium has a strong resistant layer of oxide on its surface which lowers its reactivity. (This is just like the aluminium case that you are probably familiar with.) If you add that to the trends explained below, beryllium turns out to be very unreactive.

The reactions with oxygen

Formation of simple oxides On the whole, the metals burn in oxygen to

form a simple metal oxide. Beryllium is reluctant to burn unless it is in

the form of dust or powder. Beryllium has a very strong (but very thin) layer of beryllium oxide on its surface, and this prevents any new oxygen getting at the underlying beryllium to react with it.

Thermal stability of group2 carbonates

• All the carbonates in this Group undergo thermal decomposition to give the metal oxide and carbon dioxide. Thermal decomposition is the term given to splitting up a compound by heating it.

• If "X" represents any one of the elements:

• As you go down the Group, the carbonates have to be heated more strongly before they will decompose.

• The carbonates become more stable to heat as you go down the Group.

Thermal stability of carbonates due to polarisability

• The positive ion attracts the delocalised electrons in the carbonate ion towards itself. The carbonate ion becomes polarised.

• If this is heated, the carbon dioxide breaks free to leave the metal oxide.

• How much you need to heat the carbonate before that happens depends on how polarised the ion was. if it is highly polarised, you need less heat than if it is only slightly polarised.

• The smaller the positive ion is, the higher the charge density, and the greater effect it will have on the carbonate ion. As the positive ions get bigger as you go down the Group, they have less effect on the carbonate ions near them. To compensate for that, you have to heat the compound more in order to persuade the carbon dioxide to break free and leave the metal oxide.

• In other words, as you go down the Group, the carbonates become more thermally stable.

Thermal stability of Group2 nitrates

• 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:

• As you go down the Group, the nitrates also have to be heated more strongly before they will decompose.

• The nitrates also become more stable to heat as you go down the Group.

Why the thermal stability of the metals nitrates increases down the group?

• The argument is exactly the same here as that in carbonates. The small positive ions at the top of the Group polarise the nitrate ions more than the larger positive ions at the bottom

The reason, once more, is that the polarising power of the X2+ decreases as ionic radius increases.

Solubility of group II sulphates

The solubility of these sulphates decreases as we descend the group, with barium sulphate being insoluble in water.

Two factors are involved in dissolving: • Lattice energy • Enthalpy of hydration.

• Due to the large size of the sulphate anion there is little difference between the lattice energies for these compounds. However, due to the change in ionic radius (i.e. charge density), there is significant difference in terms of their ability to hydrate.

• The greater the charge density the easier it is for the cation to hydrate and hence dissolve in water due to greater attraction with the polar water molecules.

• Magnesium sulphate dissolves in water whereas barium sulphate does not.

The relationship between enthalpy of solution and solubility

• The assumption is made that the more endothermic (or less exothermic) the enthalpy of solution is, the less soluble the compound.

• So sulphates and carbonates become less soluble as you go down the Group; hydroxides become more soluble.

Reactions of group 1 and group 2 oxides with cold water.

• All of the oxides react to give the hydroxide; peroxides and superoxides give other products as well. Group 1 hydroxides are water-soluble, group 2 oxides are sparingly soluble apart from barium hydroxide which is fairly soluble.

            MgO(s) +H2O(l)      Mg(OH)2(aq)

• All the group 2 oxides react similarly; the hydroxides are sparingly soluble and give white suspensions except for barium hydroxide. Calcium oxide reacts very exothermically.

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