The s-Block Elements • Topic 1 of 3

Group 1: Alkali Metals

The Group 1 elements — lithium, sodium, potassium, rubidium, caesium and the radioactive francium — are the alkali metals. The name comes from the fact that their hydroxides are strong alkalis. Every one of them has a single, loosely held electron in its outermost shell, giving the general valence configuration ns1. That lone electron explains almost everything about their chemistry: they lose it readily to form +1 ions, they are extremely reactive, and they are never found free in nature.

General characteristics

Alkali metals are soft (Na can be cut with a knife), silvery-white, light metals with low melting and boiling points and low densities — lithium, sodium and potassium are actually lighter than water. Because the single valence electron is shared in a weak metallic bond, the metallic bonding is feeble, which is why they are soft and melt easily.

Trends down the group

Atomic and ionic radii increase down the group as new electron shells are added; they are the largest atoms in their respective periods. Ionisation enthalpy is the lowest in each period and decreases down the group, because the valence electron is held more loosely as the atom grows and shielding increases. Hydration enthalpy decreases from Li+ to Cs+ because the smaller an ion, the more strongly it attracts water molecules — so Li+, despite being a small atom in size, is the most heavily hydrated and behaves like a large ion in solution.

Reactivity and reducing power: chemical reactivity increases down the group as ionisation enthalpy falls. All react vigorously with water to give the hydroxide and hydrogen: 2Na + 2H2O → 2NaOH + H2. In aqueous solution lithium is actually the strongest reducing agent (most negative E° = −3.04 V) because of its very high hydration enthalpy, even though caesium loses its electron most easily in the gas phase.

Flame colours: the loosely held electron is easily excited by a flame and emits visible light on returning to the ground state — Li gives crimson red, Na golden yellow, K lilac (pale violet), Rb red-violet and Cs blue. These flame tests are a classic way to identify the metal.

Anomalous behaviour of lithium

Lithium, the first member, differs from the rest because of its very small size and high polarising power (high charge density). Its compounds are more covalent (e.g. LiCl is somewhat covalent and soluble in organic solvents), Li2CO3, LiOH, LiF and Li3PO4 are sparingly soluble, and lithium reacts directly with nitrogen to form Li3N — the only alkali metal to do so. On burning in air Li forms mainly the oxide Li2O, whereas Na forms the peroxide Na2O2 and K, Rb, Cs form superoxides (KO2).

Diagonal relationship: Li and Mg

Lithium resembles magnesium (Group 2, the element diagonally down-right) because their charge-to-radius ratios are similar. Both form mainly normal oxides, both nitrides (Li3N, Mg3N2), both have sparingly soluble carbonates and fluorides, and both carbonates decompose on heating to give the oxide and CO2.

Important compounds of sodium

Sodium hydroxide (NaOH, caustic soda) is made by the electrolysis of brine (Castner–Kellner cell); it is used in soap, paper, rayon and to make many sodium salts. Sodium chloride (NaCl), common salt, is obtained from sea water and rock salt and is the raw material for NaOH, Na2CO3 and Cl2. Sodium carbonate (Na2CO3·10H2O, washing soda) is manufactured by the Solvay (ammonia-soda) process and is used in glass, soap and water softening. Sodium hydrogencarbonate (NaHCO3, baking soda) is made by passing CO2 through brine saturated with ammonia; it is a mild antacid and the source of CO2 in baking and fire extinguishers.

Characteristic flame colours of alkali and alkaline earth metalsFlame Test ColoursLicrimsonNayellowKlilacRbred-violetCsblueCabrick-redBaapple-greenThe loosely held valence electron is excited and emits a characteristic colour.
Solved Examples
1
Worked Example
Why does ionisation enthalpy decrease down Group 1 from Li to Cs?
Solution
  1. Ionisation enthalpy is the energy needed to remove the outermost ns1 electron.
  2. Down the group a new electron shell is added at each step, so atomic size increases.
  3. The valence electron is therefore farther from the nucleus, and inner shells shield it more from the nuclear charge.
  4. A more loosely held electron needs less energy to remove, so ionisation enthalpy falls steadily.

Answer: Increasing atomic size and greater shielding down the group make the valence electron easier to remove, so ionisation enthalpy decreases from Li to Cs.

2
Worked Example
Although caesium loses its electron most easily, lithium is the strongest reducing agent in aqueous solution. Explain.
Solution
  1. In aqueous solution the standard electrode potential E° depends on three steps: sublimation, ionisation and hydration of the ion.
  2. The Li+ ion is very small, so its hydration enthalpy is exceptionally large (very exothermic).
  3. This large release of hydration energy more than compensates for lithium's high ionisation enthalpy.
  4. The net effect gives lithium the most negative E° (−3.04 V), making it the strongest reducing agent in water.

Answer: Lithium's very high hydration enthalpy makes its E° the most negative, so it is the strongest reducing agent in aqueous solution despite its high ionisation enthalpy.

3
Worked Example
Identify the alkali metals giving (a) a golden-yellow flame and (b) a crimson-red flame, and explain why a flame colour appears at all.
Solution
  1. A flame supplies enough energy to excite the loosely held valence electron to a higher energy level.
  2. When the electron falls back to its ground state it emits the absorbed energy as visible light of a definite wavelength.
  3. Sodium emits at the wavelength seen as golden-yellow.
  4. Lithium emits at the wavelength seen as crimson-red.

Answer: (a) Sodium gives a golden-yellow flame; (b) lithium gives a crimson-red flame. The colour arises from emission of light as the excited valence electron returns to its ground state.

4
Worked Example
List three ways in which lithium shows anomalous behaviour compared with the other alkali metals.
Solution
  1. Lithium is the smallest alkali metal with the highest polarising power, so its compounds have appreciable covalent character (e.g. LiCl dissolves in organic solvents).
  2. Several lithium salts — Li2CO3, LiF, LiOH and Li3PO4 — are sparingly soluble in water, unlike those of Na or K.
  3. Lithium reacts directly with nitrogen to form lithium nitride, Li3N, which the other alkali metals do not do.

Answer: Covalent character of its compounds, low solubility of Li2CO3/LiF/Li3PO4, and direct formation of Li3N are three anomalies of lithium.

5
Worked Example
Write the balanced equation for the Solvay process step that produces sodium hydrogencarbonate, and name the compound commonly called washing soda.
Solution
  1. In the Solvay process, brine is first saturated with ammonia and then carbon dioxide is passed in.
  2. The reaction is: NaCl + NH3 + CO2 + H2O → NaHCO3 + NH4Cl.
  3. The sparingly soluble NaHCO3 precipitates out and is filtered off.
  4. On heating, 2NaHCO3 → Na2CO3 + H2O + CO2; the hydrated Na2CO3·10H2O is washing soda.

Answer: NaCl + NH3 + CO2 + H2O → NaHCO3 + NH4Cl; washing soda is Na2CO3·10H2O.

6
Worked Example
Arrange Li+, Na+, K+ and Cs+ in order of decreasing hydrated radius in aqueous solution, and explain.
Solution
  1. The smaller the bare ion, the higher its charge density and the more water molecules it binds tightly.
  2. Li+ is the smallest ion, so it is the most heavily hydrated and has the largest hydrated radius.
  3. As the bare ionic radius increases (Li+ < Na+ < K+ < Cs+), the degree of hydration falls.
  4. Hence the hydrated radius decreases in the same order as the bare radius increases.

Answer: Decreasing hydrated radius: Li+ > Na+ > K+ > Cs+. The smallest bare ion (Li+) attracts water most strongly and is the most hydrated.

Key Points

  • Alkali metals have the valence configuration ns1; they are soft, light, low-melting and form +1 ions.
  • Down the group atomic/ionic radii increase while ionisation enthalpy and hydration enthalpy decrease; reactivity increases.
  • Li+ is the most hydrated ion and lithium is the strongest reducing agent in water (E° = −3.04 V).
  • Flame colours: Li crimson, Na golden-yellow, K lilac, Rb red-violet, Cs blue; lithium is anomalous (covalent salts, forms Li3N) and resembles Mg diagonally.
  • Key compounds: NaOH (caustic soda), NaCl (common salt), Na2CO3·10H2O (washing soda, Solvay process), NaHCO3 (baking soda).
Quick Check — 4 MCQs
Tap an option to check your answer0 / 4
Q1.Which alkali metal gives a golden-yellow colour in a flame test?
Explanation: Sodium's excited valence electron emits light seen as golden-yellow, the basis of the sodium flame test.
Q2.Down Group 1 from Li to Cs, the ionisation enthalpy:
Explanation: Increasing atomic size and shielding make the ns1 electron easier to remove, so ionisation enthalpy decreases down the group.
Q3.Lithium resembles which Group 2 element due to a diagonal relationship?
Explanation: Li and Mg have similar charge-to-radius ratios, so they share many properties (nitride formation, sparingly soluble carbonates).
Q4.The chemical name of washing soda is:
Explanation: Washing soda is hydrated sodium carbonate, Na2CO3·10H2O, made by the Solvay process.
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