Chemical Families - Alkaline Metals and Alkaline-Earth Metals

The physical and chemical properties of elements and their compounds are highly dependent upon the number of electrons that are found in their outermost energy level.

When elements are arranged in the periodic table in order of increasing atomic number, a regular change in the outermost electron arrangement is observed. In turn a regular variation of properties is also observed.

Group 1 elements are metals. They are called alkali metals and are the most reactive group of metals. They are found on the left hand side of the periodic table.
They include Lithium, Sodium, and Potassium

Alkali metals have very similar physical and chemical properties because all alkali metals have one electron in the outermost energy level.

Group I elements show typical metal properties such as; good conductivity of heat, good conductivity of electricity, high boiling points. and shiny surface when freshly cut.

Other physical properties of alkali metals include; Low melting points, very low density (they float on water), very soft and can be cut with an ordinary knife easily.

IMPORTANT FACTS

1. The melting points and boiling points generally decrease.
2. The density generally increases.

The atomic radius is the distance between the nucleus and the outermost energy level in an atom. It is sometimes called atomic size.

Of the three group 1 elements, namely, lithium, sodium and potassium, lithium has the smallest atomic radius followed by sodium. Potassium has the largest atomic radius. The atomic radius increases as we go down the group.

This is because as we go down the group, the number of energy levels increase.

Ionic radius is the distance between the nucleus and the outermost energy level in an ion. Generally, the ionic radii increase as we go down the group.

Down the group, the number of energy levels increase which increase the ionic radius

Ionisation energy is a measure of how difficult it is to remove an electron from an atom in gaseous state. The ionisation energy is that energy that must be absorbed to remove the outermost electron from an atom.

The 1st ionisation energy for group I elements is as follows; Lithium 520kJ, sodium 496kJ and Potassium 419kJ. Notice that the 1st ionisation energy decreases as you move down the group in the periodic table. This means that it is easier to remove an electron from the atom down in Group I

When alkali metals are freshly cut, they look shiny and silvery but they tarnish immediately when exposed to air. Group 1 elements react with air to form oxides and peroxides depending on the amount of oxygen available.

Group I elements also burn with characteristic colours in a Bunsen flame. Lithium burns with a scarlet flame; sodium burns with a yellow flame; and potassium burns with a lilac flame.

Reaction of Alkali Metals with Water

A. Lithium

Lithium reacts slowly with water. It does not melt. The reaction produces heat.The heat produced is not sufficient to melt the lithium metal.
Hydrogen gas and lithium hydroxide are formed when lithium reacts with water. Since lithium hydroxide is an alkaline solution, the indictor will change colour to show the presence of an alkaline solution.

Lithium + Water → Lithium hydroxide + Hydrogen
2Li(s) + 2H₂O(l) → 2LiOH(aq) + H₂(g)

B. Sodium

Sodium reacts more vigorously than lithium. Immediately it is dropped in water, it moves around the surface of the water.
When the sodium is placed on top of a floating piece of filter paper, it catches fire.
Sodium reacts vigorously with water to form hydrogen and sodium hydroxide which is an alkaline solution. Sodium placed on the filter paper ignites because the heat produced is sufficient to ignite the hydrogen gas liberated.

Sodium + Water → Sodium hydroxide + Hydrogen
2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

C. Potassium

Potassium reacts explosively with water.Immediately it is dropped in water it ignites spontaneously. The heat produced is so much that it ignites the hydrogen produced. Potassium hydroxide solution, which is an alkaline solution, is also formed.

Potassium + Water → Potassium hydroxide + Hydrogen
2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g)

Reaction of Alkali Metals with Chlorine

Alkali metals react readily with chlorine to form chlorides. Lithium forms a white solid of lithium chloride and sodium forms a white solid of sodium chloride. Potassium reacts explosively with chlorine to form a white solid of potassium chloride.

Sodium + Chlorine → Sodium Chloride
2Na(s) + 2Cl₂O(l) → 2NaCl(aq)

Uses of Alkali Metals

1. Sodium metal is used in the preparation of tetraethyl lead, an important antiknock reagent in petrol. However, tetraethyl lead is being phased out in many countries because of lead pollution problems.
2. Sodium metal is used in the preparation of titanium metal from titanium chloride.
3. Sodium vapour is used in lamps for street lighting.
4. The sodium compound, NaCl is widely used to season or add taste to food. It is not healthy to eat too much salt.
5. Sodium carbonate is used in the manufacture of glass, detergents and for softening hard water.
6. Sodium hydroxide and sodium chloride are used in the manufacture of soaps and detergents.

The Group II elements are known collectively as alkaline-earth metals. They are found at the left hand side of the periodic table just after Group I.
Elements in this group include beryllium,magnesium and calcium.

Group II elements have two electrons in their outermost energy levels. They have similar or almost similar physical properties and chemical properties. Beryllium is the only member of this group that has unique characteristics.

Typical Metal Properties of Alkaline-earth Metals

• High melting points.
• High boiling points.
• Grey silvery surface.
• Good conductors of heat.
• Good conductors of electricity.

Important Trends down the Group II

• Melting point and boiling points generally decrease.
• Atoms get bigger.
• Ions get bigger
• Density increases.

Atomic Size and Ionic Size of Alkaline-earth Metals

The atomic size of alkaline-earth metals increases down the group.
This is because down the group, the number of energy levels increase. Another reason is that the electrons in the inner energy levels shield the electrons in the outer energy levels from full attraction by the nucleus

The ionic size of the same elements also increases down the group because of additional energy levels. The atomic radii of the elements are bigger than the ionic radii. Alkaline-earth metals form ions by losing all the electrons in the outermost energy level. The ions end up with one energy level less than that of the atom.

Ionization Energy of Alkaline-earth Metals

Ionization energy of alkaline-earth metals decreases down the group because as one goes down the group the size of the atoms increase. Because of this,the attraction of electrons in the outermost energy level by the nucleus decrease and hence the ionization energy decreases.

Generally, alkaline-earth metals burn in air to form simple metal oxides.

1. Beryllium

Beryllium is a silvery metal. It has a very strong but very thin layer of beryllium oxide on its surface which prevents any further attack from air. In fact getting to the underlying beryllium to react with air is very difficult even at over 600^oC. However, powdered beryllium metal does burn in air to give a white beryllium oxide.

Beryllium + Oxygen → Beryllium oxide
2Be(s) + 2O₂(l) → 2BeO(aq)

2. Magnesium

Magnesium burns readily with an intense bright white flame to produce a white powder of magnesium oxide.

Magnesium + Oxygen → Magnesium oxide
2Mg(s) + 2O₂(l) → 2MgO(aq)

3. Calcium

Calcium is a silvery metal. The surface of the calcium metal is covered with a thin layer of oxide that protects the metal from further attack by air. At first it is reluctant to burn because of the oxide coating, but then bursts drastically into a white flame, which burns intensely to form a white solid of calcium oxide

Calcium + Oxygen → Calcium oxide
2Ca(s) + 2O₂(l) → 2CaO(aq)

NOTE: Beryllium, magnesium and calcium do not form peroxides when heated in air. The reactivity increases as you go down the group due to increasing ease of removing the electrons in the outermost energy level.

Reaction of Alkaline-earth metals (Mg and Ca) with Water

Beryllium does not react with water or steam even when it is red-hot.

Very clean magnesium slightly reacts with cold water to form magnesium hydroxide solution which turns red litmus blue and hydrogen gas

Magnesium + water → magnesium hydroxide + Hydrogen
Mg(s) + 2H₂O(l) → Mg(OH)₂(aq) + H₂(g)

Calcium reacts with cold water with increased vigour than magnesium to give calcium hydroxide solution and hydrogen gas.

Calcium + water → Calcium hydroxide + Hydrogen
Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g)

NOTE: The reactivity of alkaline-earth metals with water shows increasing reactivity as you go down the group. This is mainly due to decreasing ionization energy down the group.

Reaction of Alkaline-earth metals (Mg and Ca) with Chlorine

Alkaline-earth metals react with chlorine to form chloride salts.
Heated beryllium metal reacts with chlorine to form beryllium chloride.

Beryllium + Chlorine → Beryllium chloride
Be(s) + Cl₂(l) → BeCl₂(aq)

Magnesium burns in chlorine with a bright light to form a white solid of magnesium chloride.

Magnesium + Chlorine → Magnesium chloride
Mg(s) + Cl₂(l) → MgCl₂(aq)

Calcium appears to burn slowly in chlorine compared to magnesium to form a white solid of calcium chloride.

Calcium + Chlorine → Calcium chloride
Ca(s) + Cl₂(l) → CaCl₂(aq)

Reactivity of alkaline-earth metals with chlorine increases as you go down the group.
The calcium oxide coating formed on the surface as a result of its reactivity with oxygen forms a protective coat that make the reaction of calcium to appear slower. The inner silvery calcium metal cannot be easily reached by the chlorine.

Reaction of alkaline-earth metals with dilute acids

Reaction of calcium and acids is very explosive. Very dilute acids should be used. Alkaline-earth metals react with dilute acids to form a salt and hydrogen only. The evolution of hydrogen gas is evident by the 'pop' sound produced by the burning splint. Very dilute nitric acid should also be used with magnesium.

Uses of Alkaline-earth Metals

(a) Beryllium

• Used in transmission of X-rays (beryllium transmits X-rays better than aluminium).
• Beryllium alloyed with copper gives a hard strong alloy with high resistance to wear. It is therefore used in computer parts, and other instruments with desirable lightness and stiffness.
• Alloys of beryllium are used as structural materials for high performance aircrafts, missiles, spacecraft and communication satellites among many other things.
• The oxide is used in the nuclear industry.

(b) Magnesium

• It is lighter than aluminium, and it is used to make alloys used for aircraft, car engine casings, and missile construction.
• It is used as a reducing agent for the production of uranium and other metals from their salts.
• Magnesium hydroxide (milk of magnesia), chloride, sulphate (Epsom salts), among others are used in medicine.
• Magnesium oxide is used as brick-liners in furnaces.
• Used in computers for radio-frequency shielding.

(c) Calcium

• It is used as a reducing agent in the preparation of metals such as thorium, uranium, zirconium, etc.
• Calcium forms calcium carbonate which is a component of Portland cement.
• Calcium carbonate is used as antacid tablets.