Why Sodium Is Stored Under Kerosene

Sodium is soft enough to cut with a knife, reacts with water fast enough to catch fire, and must be kept under kerosene. Its compounds, however, are among the most useful in any home or factory: caustic soda makes soap and paper, washing soda softens hard water, and baking soda makes batter rise.

This lesson covers the general trends of Group 1 elements, and the preparation, properties and uses of sodium hydroxide, sodium carbonate and sodium bicarbonate.

**The Group 1 alkali metals — lithium, sodium, potassium, rubidium and caesium — all have an $n{s}^1$ configuration, so they readily lose one electron, and down the group their atoms grow larger, their ionisation enthalpy falls and their reactivity rises.

Physical trends down the group:

- Atomic and ionic radii increase as new shells are added
- Ionisation enthalpy** decreases: lithium 520, sodium 496, potassium 419, rubidium 403 and caesium 376 kJ mol${}^{-1}$
- Melting and boiling points decrease, because metallic bonding weakens as the atoms grow
- Density generally increases, though potassium is lighter than sodium
- Flame colours: lithium crimson red, sodium golden yellow, potassium lilac, rubidium red-violet and caesium blue, because the loosely held electron is easily excited

Chemical trends:

- Reaction with water grows more violent down the group: $2M+2{\mathrm{H}}_2\mathrm{O}\to 2MOH+{\mathrm{H}}_2$
- Reaction with oxygen: lithium forms the oxide $\mathrm{L}{\mathrm{i}}_2\mathrm{O}$, sodium the peroxide $\mathrm{N}{\mathrm{a}}_2{\mathrm{O}}_2$, and potassium the superoxide $\mathrm{K}{\mathrm{O}}_2$
- Solutions in liquid ammonia are deep blue, because of electrons surrounded by ammonia molecules

Anomalous lithium. Its tiny size and high charge density make lithium unlike the rest of the group and similar to magnesium, its diagonal neighbour: it forms a nitride, $\mathrm{L}{\mathrm{i}}_3\mathrm{N}$, and its carbonate decomposes on heating.

Worked reasoning — why lithium is the strongest reducing agent in water. Lithium has the highest ionisation enthalpy in the group, yet the most negative electrode potential, $E^{\circ }=-3.04$ V. The tiny $\mathrm{L}{\mathrm{i}}^{+}$ ion releases a very large hydration enthalpy when water molecules surround it, which more than repays the energy needed to remove its electron.

An everyday example. Sodium vapour street lamps along Indian highways glow golden yellow for the same reason sodium colours a flame: excited electrons in sodium atoms fall back and emit yellow light.

The substance. Sodium is stored under kerosene because it reacts quickly with the oxygen and moisture in air, and kerosene contains neither.

Sodium hydroxide is made by electrolysing brine, sodium carbonate by the Solvay ammonia-soda process, and sodium bicarbonate as an intermediate of that process — and their uses range from soap and glass to baking and antacids.

Sodium hydroxide, NaOH (caustic soda):

- Preparation by electrolysis of brine, in the Castner-Kellner cell with a mercury cathode or in a modern membrane cell:

$$2NaCl+2{\mathrm{H}}_2\mathrm{O}\to 2NaOH+\mathrm{C}{\mathrm{l}}_2+{\mathrm{H}}_2$$

- Properties: white, deliquescent solid; strongly basic; absorbs carbon dioxide from air, $2NaOH+\mathrm{C}{\mathrm{O}}_2\to \mathrm{N}{\mathrm{a}}_2\mathrm{C}{\mathrm{O}}_3+{\mathrm{H}}_2\mathrm{O}$
- Uses: making soap, paper, artificial silk and dyes; refining petroleum; purifying bauxite

**Sodium carbonate, $\mathrm{N}{\mathrm{a}}_2\mathrm{C}{\mathrm{O}}_3\cdot 10{\mathrm{H}}_2\mathrm{O}$ (washing soda):

- Preparation by the Solvay process** — brine saturated with ammonia is treated with carbon dioxide:

$$\mathrm{NaCl}+\mathrm{N}{\mathrm{H}}_3+\mathrm{C}{\mathrm{O}}_2+{\mathrm{H}}_2\mathrm{O}\to \mathrm{NaHC}{\mathrm{O}}_3+\mathrm{N}{\mathrm{H}}_4\mathrm{Cl}$$

$$2NaHC{\mathrm{O}}_3\to \mathrm{N}{\mathrm{a}}_2\mathrm{C}{\mathrm{O}}_3+\mathrm{C}{\mathrm{O}}_2+{\mathrm{H}}_2\mathrm{O}$$

- Sparingly soluble sodium bicarbonate crystallises out, and heating it gives soda ash
- Ammonia is recovered by heating ammonium chloride with slaked lime: $2N{\mathrm{H}}_4\mathrm{Cl}+\mathrm{Ca}(\mathrm{OH}{)}_2\to 2N{\mathrm{H}}_3+\mathrm{CaC}{\mathrm{l}}_2+2{\mathrm{H}}_2\mathrm{O}$
- Properties: the crystals lose water to air (efflorescence); strong heating gives anhydrous soda ash; its solution is alkaline by hydrolysis
- Uses: softening hard water, making glass and soap, and laundry washing

**Sodium bicarbonate, $\mathrm{NaHC}{\mathrm{O}}_3$ (baking soda):

- Preparation**: saturating sodium carbonate solution with carbon dioxide, $\mathrm{N}{\mathrm{a}}_2\mathrm{C}{\mathrm{O}}_3+{\mathrm{H}}_2\mathrm{O}+\mathrm{C}{\mathrm{O}}_2\to 2NaHC{\mathrm{O}}_3$
- Properties: mildly alkaline; decomposes on heating to release carbon dioxide
- Uses: baking powder, antacids, a mild antiseptic for skin, and fire extinguishers

Worked example. From $2NaHC{\mathrm{O}}_3\to \mathrm{N}{\mathrm{a}}_2\mathrm{C}{\mathrm{O}}_3+\mathrm{C}{\mathrm{O}}_2+{\mathrm{H}}_2\mathrm{O}$, 168 g of baking soda releases 22.4 L of carbon dioxide at STP, so 84 g releases 11.2 L.

An everyday example. Dhokla batter in a Gujarati kitchen puffs up because baking soda releases bubbles of carbon dioxide when mixed with an acid such as lemon juice and steamed — while soda ash plants on the Gujarat coast run the Solvay process close to supplies of salt and limestone.

The substance. The Solvay process cannot make potassium carbonate — potassium bicarbonate is too soluble to crystallise out of the solution.