Reactions of the Alkali Metals, and Their Oxides and Hydroxides – Trends in Chemical Properties of Group 1 Elements and Their Compounds (Senior 4 Chemistry) | Ganzaa

Trends in Chemical Properties of Group 1 Elements and Their Compounds

Lessons

Senior 4 · Chemistry

Trends in Chemical Properties of Group 1 Elements and Their Compounds

4 lessons

Lesson 2 of 4

Reactions of the Alkali Metals, and Their Oxides and Hydroxides

25 min

Faster and faster

Drop a small piece of lithium into water and it fizzes steadily. Sodium melts into a silver ball and skates over the surface. Potassium bursts into a lilac flame. Rubidium and caesium explode. The same reaction is happening each time, but faster at every step down the group. In this lesson you will write the equations for the reactions of the alkali metals with oxygen, water and the halogens, explain the trend from their structure, and meet the compounds they make: oxides, peroxides, superoxides and the strong alkalis, the hydroxides.

Why reactivity increases down the group

In every reaction an alkali metal atom loses its outer electron: $\mathrm{M} \rightarrow \mathrm{M^{+}} + e^-$ . Down the group the outer electron is further from the nucleus and more shielded, so the first ionisation energy falls (Li 520 to Cs 376 kJ mol⁻¹). The electron is lost more easily, so the metal reacts faster. The alkali metals are strong reducing agents: they give electrons to other substances.

Reaction with oxygen

A freshly cut surface tarnishes in seconds in air. Heated in oxygen, the metals burn with coloured flames (lithium red, sodium yellow-orange, potassium lilac) to form white or pale yellow solids. The product changes down the group:

Metal	Main product in excess oxygen	Anion	Equation
lithium	lithium oxide	oxide, $\mathrm{O^{2-}}$	$4\mathrm{Li(s)} + \mathrm{O_2(g)} \rightarrow 2\mathrm{Li_2O(s)}$
sodium	sodium peroxide	peroxide, $\mathrm{O_2^{2-}}$	$2\mathrm{Na(s)} + \mathrm{O_2(g)} \rightarrow \mathrm{Na_2O_2(s)}$
potassium, rubidium, caesium	superoxide	superoxide, $\mathrm{O_2^{-}}$	$\mathrm{K(s)} + \mathrm{O_2(g)} \rightarrow \mathrm{KO_2(s)}$

Why the change? The peroxide and superoxide ions are large. The small $\mathrm{Li^{+}}$ ion has a high charge density and distorts (polarises) a large anion so much that it breaks down to the simple oxide. The larger $\mathrm{K^{+}}$ , $\mathrm{Rb^{+}}$ and $\mathrm{Cs^{+}}$ ions polarise less and pack better with a large anion, so the superoxides are stable. With a limited supply of air, sodium also forms the normal oxide: $4\mathrm{Na} + \mathrm{O_2} \rightarrow 2\mathrm{Na_2O}$ .

Definition

Peroxide ion, $\mathrm{O_2^{2-}}$ : two oxygen atoms joined by a single bond, each with oxidation state −1. Superoxide ion, $\mathrm{O_2^{-}}$ : two oxygen atoms with one extra electron between them, oxidation state −½ each. In the normal oxide ion, $\mathrm{O^{2-}}$ , oxygen is −2.

Reaction with water

Every alkali metal reacts with cold water to give the metal hydroxide and hydrogen:

2\mathrm{M(s)} + 2\mathrm{H_2O(l)} \longrightarrow 2\mathrm{MOH(aq)} + \mathrm{H_2(g)}

Table 2.1: What you see

Metal	Observations
lithium	floats; fizzes steadily; slowly disappears; the solution becomes alkaline
sodium	floats; melts into a shiny ball (the reaction gives out enough heat to melt it, at 98 °C); moves rapidly over the surface; fizzes; may burn with a yellow flame
potassium	floats; melts; the hydrogen ignites at once with a lilac flame; crackles and may spit
rubidium, caesium	sink (denser than water); react explosively; never used in schools

Potassium reacting violently with water. The hydrogen given off ignites with a lilac flame, coloured by potassium vapour, and the phenolphthalein in the water turns violet-red as potassium hydroxide forms. Photo: Ozone aurora (Philip Evans), CC BY-SA 3.0, via Wikimedia Commons

Worked example 1: sodium and water

A teacher adds 0.46 g of sodium to 200 cm³ of water. Calculate (a) the volume of hydrogen formed at room temperature and pressure (molar volume 24.0 dm³ mol⁻¹), (b) the concentration of sodium hydroxide in the solution, assuming the volume stays 200 cm³. (Na = 23.0)

Solution. $n(\mathrm{Na}) = \frac{0.46}{23.0} = 0.020\,\mathrm{mol}$ . From $2\mathrm{Na} + 2\mathrm{H_2O} \rightarrow 2\mathrm{NaOH} + \mathrm{H_2}$ :

(a) $n(\mathrm{H_2}) = \frac{0.020}{2} = 0.010\,\mathrm{mol}$ ; volume $= 0.010 \times 24.0 = 0.24\,\mathrm{dm^3} = 240\,\mathrm{cm^3}$ .

(b) $n(\mathrm{NaOH}) = 0.020\,\mathrm{mol}$ ; concentration $= \frac{0.020}{0.200} = 0.10\,\mathrm{mol\,dm^{-3}}$ .

Answer: 240 cm³ of hydrogen and a 0.10 mol dm⁻³ solution of sodium hydroxide.

Reaction with the halogens

Heated alkali metals burn in chlorine, bromine vapour or iodine vapour to form white ionic halides:

2\mathrm{Na(s)} + \mathrm{Cl_2(g)} \longrightarrow 2\mathrm{NaCl(s)} \qquad 2\mathrm{K(s)} + \mathrm{Br_2(g)} \longrightarrow 2\mathrm{KBr(s)}

The reactivity increases down Group 1 and decreases down Group 17, so the most vigorous combinations are caesium with fluorine. All the alkali metal halides are white crystalline solids with high melting points, and all are soluble in water except lithium fluoride, which is only slightly soluble.

Oxides and hydroxides: strong bases

All the Group 1 oxides are basic: they react with water to give hydroxides, and with acids to give salts.

\mathrm{Li_2O(s)} + \mathrm{H_2O(l)} \rightarrow 2\mathrm{LiOH(aq)} \qquad \mathrm{Na_2O_2(s)} + 2\mathrm{H_2O(l)} \rightarrow 2\mathrm{NaOH(aq)} + \mathrm{H_2O_2(aq)}

2\mathrm{KO_2(s)} + 2\mathrm{H_2O(l)} \rightarrow 2\mathrm{KOH(aq)} + \mathrm{H_2O_2(aq)} + \mathrm{O_2(g)}

The hydroxides are white solids that dissolve in water to give strong alkalis, fully dissociated into $\mathrm{M^{+}}$ and $\mathrm{OH^{-}}$ ions. Their solubility, and their strength as bases in the solid state, increase down the group: lithium hydroxide is the least soluble. They neutralise acids and absorb carbon dioxide:

\mathrm{NaOH(aq)} + \mathrm{HCl(aq)} \rightarrow \mathrm{NaCl(aq)} + \mathrm{H_2O(l)} \qquad 2\mathrm{NaOH(aq)} + \mathrm{CO_2(g)} \rightarrow \mathrm{Na_2CO_3(aq)} + \mathrm{H_2O(l)}

Worked example 2: oxidation states and products

Give the oxidation state of oxygen in $\mathrm{Li_2O}$ , $\mathrm{Na_2O_2}$ and $\mathrm{KO_2}$ . Hence explain why $\mathrm{KO_2}$ gives off oxygen gas with water but $\mathrm{Li_2O}$ does not.

Solution. Each alkali metal is +1. $\mathrm{Li_2O}$ : $2 (+ 1) + x = 0$ , $x = - 2$ . $\mathrm{Na_2O_2}$ : $2 (+ 1) + 2 x = 0$ , $x = - 1$ . $\mathrm{KO_2}$ : $+ 1 + 2 x = 0$ , $x = -\frac{1}{2}$ .

In $\mathrm{Li_2O}$ oxygen is already at −2, the same as in water and hydroxide, so it simply forms LiOH. In $\mathrm{KO_2}$ oxygen is at −½: with water some of it is reduced (to −1 in $\mathrm{H_2O_2}$ ) and some oxidised (to 0 in $\mathrm{O_2}$ ), so oxygen gas is released.

Warning

Common mistakes. (1) Writing NaO or $\mathrm{KO}$ : the oxide is $\mathrm{Na_2O}$ , and $\mathrm{KO_2}$ is a superoxide containing $\mathrm{O_2^{-}}$ . (2) Forgetting to balance the hydrogen: $2\mathrm{M} + 2\mathrm{H_2O} \rightarrow 2\mathrm{MOH} + \mathrm{H_2}$ . (3) Saying lithium is most reactive because it has the highest ionisation energy: it is the least reactive.

Safety. The water reactions are teacher demonstrations: rice-grain-sized pieces only, a large trough, a safety screen, goggles for everyone, and no more than lithium and sodium (potassium only by an experienced teacher). Sodium hydroxide solution damages the eyes even when dilute.

Tip

For any Group 1 reaction, write the general equation with M and then substitute: $2\mathrm{M} + 2\mathrm{H_2O} \rightarrow 2\mathrm{MOH} + \mathrm{H_2}$ ; $2\mathrm{M} + \mathrm{X_2} \rightarrow 2\mathrm{MX}$ . Only the reaction with oxygen needs a different product for each metal.

OpenLearn from The Open University: Alkali metals in water (2.5 min). Watch the reactions of the alkali metals with water one after another and compare how vigorous each one is.

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Practice

Write balanced equations with state symbols for the reactions of (a) lithium with water, (b) potassium with chlorine, (c) lithium burning in oxygen.
Describe what you would see when a small piece of sodium is added to water containing universal indicator.
Explain why potassium reacts with water more vigorously than sodium.
Name the compound and identify the anion formed when (a) sodium and (b) caesium burn in excess oxygen.
Write an equation for the reaction of sodium oxide with water, and for potassium hydroxide with sulfuric acid.
0.14 g of lithium reacts completely with water. Calculate the volume of hydrogen at room temperature and pressure (24.0 dm³ mol⁻¹; Li = 6.9).
Potassium superoxide is used in breathing equipment for miners and fire-fighters: it reacts with exhaled carbon dioxide, $4\mathrm{KO_2} + 2\mathrm{CO_2} \rightarrow 2\mathrm{K_2CO_3} + 3\mathrm{O_2}$ . What volume of oxygen at RTP is released when 1.0 mol of $\mathrm{CO_2}$ is absorbed?

Show answersHide answers

(a) $2\mathrm{Li(s)} + 2\mathrm{H_2O(l)} \rightarrow 2\mathrm{LiOH(aq)} + \mathrm{H_2(g)}$ ; (b) $2\mathrm{K(s)} + \mathrm{Cl_2(g)} \rightarrow 2\mathrm{KCl(s)}$ ; (c) $4\mathrm{Li(s)} + \mathrm{O_2(g)} \rightarrow 2\mathrm{Li_2O(s)}$ .
It floats, melts into a shiny ball, moves about on the surface and fizzes; it gets smaller and disappears; the indicator turns blue/purple as the solution becomes alkaline.
Potassium's outer electron is in the 4s orbital, further from the nucleus and more shielded than sodium's 3s electron, so it is lost more easily (lower ionisation energy).
(a) Sodium peroxide, $\mathrm{Na_2O_2}$ , peroxide ion $\mathrm{O_2^{2-}}$ . (b) Caesium superoxide, $\mathrm{CsO_2}$ , superoxide ion $\mathrm{O_2^{-}}$ .
$\mathrm{Na_2O(s)} + \mathrm{H_2O(l)} \rightarrow 2\mathrm{NaOH(aq)}$ ; $2\mathrm{KOH(aq)} + \mathrm{H_2SO_4(aq)} \rightarrow \mathrm{K_2SO_4(aq)} + 2\mathrm{H_2O(l)}$ .
$n(\mathrm{Li}) = \frac{0.14}{6.9} = 0.0203\,\mathrm{mol}$ ; $n(\mathrm{H_2}) = 0.0101\,\mathrm{mol}$ ; volume $= 0.0101 \times 24.0 = 0.24\,\mathrm{dm^3}$ (243 cm³).
2 mol $\mathrm{CO_2}$ release 3 mol $\mathrm{O_2}$ , so 1.0 mol releases 1.5 mol: $1.5 \times 24.0 = 36\,\mathrm{dm^3}$ .

Quick check

What is the main product when potassium burns in excess oxygen?

Key points

Reactivity increases down Group 1 because the outer electron is lost more easily (lower ionisation energy); the metals are strong reducing agents.
Burning in oxygen: $\mathrm{Li_2O}$ (oxide), $\mathrm{Na_2O_2}$ (peroxide), $\mathrm{KO_2}$ (superoxide).
With water: $2\mathrm{M} + 2\mathrm{H_2O} \rightarrow 2\mathrm{MOH} + \mathrm{H_2}$ ; with halogens: $2\mathrm{M} + \mathrm{X_2} \rightarrow 2\mathrm{MX}$ .
The oxides are basic; the hydroxides are strong alkalis whose solubility increases down the group.

Sources and credits

Image: Potassium & water 39.jpg by Ozone aurora (Philip Evans), via Wikimedia Commons, licensed under CC BY-SA 3.0.
Video: Alkali metals in water - Chemical elements: properties and reactions (18), OpenLearn from The Open University, embedded from YouTube.
The text, diagrams, tables, worked examples and Practice items were written and drawn by Ganzaa.
This lesson adapts CC BY-SA material and is shared under CC BY-SA 4.0.

1Faster and faster

2Why reactivity increases down the group

3Reaction with oxygen

4Reaction with water

5Reaction with the halogens

6Oxides and hydroxides: strong bases