Revision:Chemistry Unit 1.6 - Groups 1 and 2

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Revision:Chemistry Unit 1.6 - Groups 1 and 2

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A Level Chemistry Unit 1.6 - Groups 1 and 2

Physical properties

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 down.

Flame colours

Group1

Lithium – red
Sodium - orange/yellow
Potassium - lilac
Rubidium – Reddish violet
Caesium - Blue/violet

Group 2

Calcium – Orange/red
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 colour.

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 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 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

Oxygen

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.

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.

Chlorine

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.

Water

Lithium -Lithium's density is only about half that of water so it floats on the surface, gently fizzing and giving off hydrogen. It gradually reacts and disappears, forming a colourless solution of lithium hydroxide. The reaction generates heat too slowly and lithium's melting point is too high for it to melt (see sodium below).

Sodium - Sodium also floats on the surface, but enough heat is given off to melt the sodium (sodium has a lower melting point than lithium and the reaction produces heat faster) and it melts almost at once to form a small silvery ball that dashes around the surface. A white trail of sodium hydroxide is seen in the water under the sodium, but this soon dissolves to give a colourless solution of sodium hydroxide. The sodium moves because it is pushed around by the hydrogen which is given off during the reaction. If the sodium becomes trapped on the side of the container, the hydrogen may catch fire to burn with an orange flame. The colour is due to contamination of the normally blue hydrogen flame with sodium compounds.

Potassium - Potassium behaves rather like sodium except that the reaction is faster and enough heat is given off to set light to the hydrogen. This time the normal hydrogen flame is contaminated by potassium compounds and so is coloured lilac (a faintly bluish pink).

Rubidium - Rubidium is denser than water and so sinks. It reacts violently and immediately, with everything spitting out of the container again. Rubidium hydroxide solution and hydrogen are formed.

Caesium - Caesium explodes on contact with water, quite possibly shattering the container. Caesium hydroxide and hydrogen are formed.

Group 2 reactions with O₂, Cl₂ and H₂O

Oxygen

All the group 2 metals burn with oxygen to form the metal oxide with formula XO, even if some don't burn very well. Strontium and barium will also form the metal peroxides with the formula $\mathsf{XO_2}$ . All the group 2 metals also form the metal nitrides when burnt in air, which have the formula $\mathsf{X_3N_2}$ .

Water

Beryllium - Beryllium has no reaction with water or steam even at red heat.
Magnesium - Magnesium burns in steam to produce magnesium oxide and hydrogen. 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 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.

Group 1 and 2 oxides with water and dilute acid

The simple oxides

Reaction with water - These are simple basic oxides, reacting with water to give the metal hydroxide.

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.

The peroxides

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. If the temperature increases, the hydrogen peroxide produced decomposes into water and oxygen.
Reaction with dilute acids - These reactions are even more exothermic than the ones with water. A solution containing a salt and hydrogen peroxide is formed. The hydrogen peroxide will decompose to give water and oxygen if the temperature rises - again, it is almost impossible to avoid this.

The superoxides

Reaction with water - This time, a solution of the metal hydroxide and hydrogen peroxide is formed, but oxygen gas is given off as well. Once again, these are strongly exothermic reactions and the heat produced will inevitably decompose the hydrogen peroxide to water and more oxygen.

Solubility of group 2 sulphates and hydroxides

The solubilities of group 2 sulphates decreases as you go down the group but the solubilities of the hydroxides increases as you go down the group.

Thermal stabilities of nitrates and carbonates

Group 1

Heating the nitrates

Most nitrates tend to decompose on heating to give the metal oxide, brown fumes of nitrogen dioxide, and oxygen. For example, a typical Group 2 nitrate like magnesium nitrate decomposes like this:

$\mathsf{2Mg(NO_3)_2 \longrightarrow 2MgO + 4NO_2 + O_2}$

In Group 1, lithium nitrate behaves in the same way - producing lithium oxide, nitrogen dioxide and oxygen:

$\mathsf{4LiNO_3 \longrightarrow 2Li_2O + 4NO_2 + O_2}$

The rest of the Group, however, don't decompose so completely (at least not at Bunsen temperatures) - producing the metal nitrite and oxygen, but no nitrogen dioxide.

$\mathsf{2XNO_3 \longrightarrow 2XNO_2 + O_2}$

All the nitrates from sodium to caesium decomposes in this same way, the only difference being how hot they have to be to undergo the reaction. As you go down the Group, the decomposition gets more difficult, and you have to use higher temperatures.

Heating the carbonates

Most carbonates tend to decompose on heating to give the metal oxide and carbon dioxide. Group 1 carbonates don't decompose at Bunsen temperatures, although at higher temperatures they will. The decomposition temperatures again increase as you go down the Group.

The thermal stability of the hydrogencarbonates

The Group 1 hydrogencarbonates are stable enough to exist as solids, although they do decompose easily on heating.

Group 2

Heating 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:

$\mathsf{XCO_3 \longrightarrow XO + CO_2}$

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

Heating 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:

$\mathsf{2X(NO_3)_2 \longrightarrow 2XO + 4NO_2 + O_2}$

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

Explanation

A small 2+/1+ ion has a lot of charge packed into a small volume of space. It has a high charge density and will have a marked distorting effect on any negative ions which happen to be near it.

A bigger 2+/1+ ion has the same charge spread over a larger volume of space. Its charge density will be lower, and it will cause less distortion to nearby negative ions.

The structure of the carbonate ion

If you worked out the structure of a carbonate ion using "dots-and-crosses" or some similar method, you would probably come up with:

(relevant diagram needed here)

This shows two single carbon-oxygen bonds and one double one, with two of the oxygens each carrying a negative charge. Unfortunately, in real carbonate ions all the bonds are identical, and the charges are spread out over the whole ion - although concentrated on the oxygen atom, the charges are delocalised.

Polarising the carbonate ion

Now imagine what happens when this ion is placed next to a positive ion. 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.

The argument for the nitrates is exactly the same. The small positive ions at the top of the Groups polarise the nitrate ions more than the larger positive ions at the bottom.