Group IA–IIA metals and water hardness

The Group IA elements (alkali metals) and Group IIA elements (alkaline-earth metals) are the most reactive metals. Their atoms give up one or two outer electrons very easily, so in nature they occur only in compounds, react vigorously with water and form alkalis (hence the name). Sodium and potassium ions carry nerve impulses, calcium builds bones and teeth, and magnesium is part of chlorophyll. Table salt, soda, lime, gypsum and marble — some of the most common substances in daily life — are compounds of these elements. Hardness of water also belongs to this topic.

§1. Alkali metals: structure and physical properties

With one outer electron and a large atomic radius, the alkali metals lose that electron very easily — they are the strongest reducing agents. Down the group the radius grows, ionisation energy falls and reactivity rises: lithium reacts gently with water, sodium vigorously, potassium bursts into flame, and rubidium and caesium explode.

Physical properties: silvery-white, very soft (they can be cut with a knife), light (lithium is the least dense metal; Li, Na and K float on water), with low melting points (caesium melts in the hand, at about 28 °C). Because they react quickly with oxygen and moisture, they are stored under paraffin oil (lithium under petroleum jelly). Their salts colour a flame: Li red, Na yellow, K lilac — a way to identify them.

§2. Chemical properties of the alkali metals

• With water: 2Na + 2H₂O → 2NaOH + H₂↑ — an alkali and hydrogen form.
• The product with oxygen depends on the metal: 4Li + O₂ → 2Li₂O (oxide); 2Na + O₂ → Na₂O₂ (peroxide); K + O₂ → KO₂ (superoxide). To make the oxide Na₂O, the peroxide is heated with sodium.
• With halogens and sulfur: 2Na + Cl₂ → 2NaCl; 2Na + S → Na₂S.
• With hydrogen they form hydrides (hydrogen −1): 2Na + H₂ → 2NaH; a hydride releases hydrogen with water: NaH + H₂O → NaOH + H₂.
• Lithium reacts even with nitrogen at room temperature: 6Li + N₂ → 2Li₃N.

Preparation: electrolysis of a molten salt or hydroxide: 2NaCl → 2Na + Cl₂.

Sodium peroxide releases oxygen with water and CO₂, so it is used to regenerate air in submarines and spacecraft:

Important compounds: NaOH (caustic soda), Na₂CO₃ (washing soda), NaHCO₃ (baking soda; on heating 2NaHCO₃ → Na₂CO₃ + CO₂ + H₂O, which makes dough rise), NaCl (table salt), and KNO₃ and KCl (fertilisers).

Worked example

4.6 g of sodium was dropped into 95.6 g of water. Find the mass fraction of NaOH in the solution.

1. n(Na) = 4.6/23 = 0.2 mol → NaOH 0.2 mol = 8 g, H₂ 0.1 mol = 0.2 g.
2. Mass of solution: 4.6 + 95.6 − 0.2 = 100 g.
3. ω = 8/100 = 0.08, i.e. 8 %.

Worked example

3.9 g of an unknown alkali metal reacted with water, releasing 1.12 L of hydrogen (STP). Identify the metal.

1. n(H₂) = 1.12/22.4 = 0.05 mol.
2. 2Me + 2H₂O → 2MeOH + H₂: n(Me) = 0.1 mol.
3. M = 3.9/0.1 = 39 g/mol — potassium.

§3. Alkaline-earth metals

They are harder and less reactive than the alkali metals, but reactivity increases down the group. Hydroxide character also changes: Be(OH)₂ is amphoteric (it reacts with acids and alkalis: Be(OH)₂ + 2NaOH → Na₂[Be(OH)₄]), Mg(OH)₂ is a weak insoluble base, and Ca(OH)₂, Sr(OH)₂ and Ba(OH)₂ are alkalis.

• Ca, Sr and Ba react with cold water: Ca + 2H₂O → Ca(OH)₂ + H₂; magnesium only with hot water; beryllium not at all.
• Magnesium burns in air with a dazzling white flame and keeps burning even in CO₂: 2Mg + CO₂ → 2MgO + C; with nitrogen it forms a nitride: 3Mg + N₂ → Mg₃N₂. So burning magnesium cannot be put out with carbon dioxide.
• Flame colours: Ca brick-red, Sr crimson, Ba apple-green.

Calcium compounds: CaO — quicklime; Ca(OH)₂ — slaked lime, whose clear solution is limewater (the test for CO₂: it turns milky); CaCO₃ — chalk, limestone, marble; CaSO₄·2H₂O — gypsum; Ca₃(PO₄)₂ — phosphorite.

§4. Hardness of water

Hard water contains a lot of Ca²⁺ and Mg²⁺ ions. Soap does not lather in it, and scale (mainly CaCO₃) forms in boilers and kettles, reducing heat transfer.

• Temporary (carbonate) hardness is caused by hydrogencarbonates Ca(HCO₃)₂ and Mg(HCO₃)₂; it is removed by boiling:

It can also be removed with lime: Ca(HCO₃)₂ + Ca(OH)₂ → 2CaCO₃↓ + 2H₂O.

• Permanent hardness is caused by sulfates and chlorides (CaSO₄, MgCl₂); boiling does not help, so washing soda or a phosphate is added:

• The universal method is ion exchange: a cation-exchange resin swaps the Ca²⁺ and Mg²⁺ in water for Na⁺ or H⁺.

Worked example

1 L of water contains 0.162 g of Ca(HCO₃)₂. What mass of Ca(OH)₂ is needed to soften it?

1. n(Ca(HCO₃)₂) = 0.162/162 = 0.001 mol.
2. The equation gives a 1 : 1 ratio → n(Ca(OH)₂) = 0.001 mol.
3. m = 0.001·74 = 0.074 g.

§5. Identification tests and uses

The Ba²⁺ ion gives a white precipitate with sulfate: Ba²⁺ + SO₄²⁻ → BaSO₄↓ (insoluble in acids); Ca²⁺ gives a white precipitate with carbonate or ethanedioate (oxalate). Salts of sodium, potassium, calcium, strontium and barium are identified by flame colour.
Uses: NaCl is a raw material for the food and chemical industries; NaOH is used to make soap and paper; Na₂CO₃ to make glass and detergents; potassium salts are fertilisers; magnesium forms light alloys (aircraft); calcium compounds are building materials (lime, cement, plaster); barium sulfate is a contrast agent for X-ray examinations; lithium is used in rechargeable batteries.

Check yourself

1. 7.8 g of potassium reacted with water. What volume of H₂ (STP) is released?

2. Which hydroxide is amphoteric: Mg(OH)₂, Be(OH)₂, Ca(OH)₂?

3. Can permanent hardness be removed by boiling?

4. 15.6 g of Na₂O₂ reacted completely with water. What volume of O₂ (STP) is released?

5. 10 g of CaCO₃ was roasted and the CaO slaked with water. What mass of Ca(OH)₂ forms?

6. How can a sodium salt be identified with a flame?

7. What mainly forms when sodium burns in oxygen?

8. Why can’t burning magnesium be put out with CO₂?

Common mistakes

• Writing that sodium gives only Na₂O in oxygen: mainly Na₂O₂ forms.
• Not subtracting the H₂ released from the mass of the solution.
• Mixing up the methods for removing temporary and permanent hardness.
• Calling Be(OH)₂ an alkali: it is amphoteric.
• Writing that alkali metals displace other metals from salt solutions: they react with the water first.