Reactions of the Period 3 Elements – Trends in Chemical Properties of Period 3 Elements and Their Compounds (Senior 4 Chemistry) | Ganzaa

Trends in Chemical Properties of Period 3 Elements and Their Compounds

Lessons

Senior 4 · Chemistry

Trends in Chemical Properties of Period 3 Elements and Their Compounds

4 lessons

Lesson 2 of 4

Reactions of the Period 3 Elements

25 min

Same water, three different answers

Drop a rice-grain of sodium into a trough of water and it melts into a silver ball that skates about, fizzing, until it vanishes. Leave a clean strip of magnesium ribbon in cold water and you may need a day to collect a test tube of gas. Boil water in an aluminium pan every morning for years and the pan does not dissolve. All three are metals from the same period, and all three can in principle reduce water to hydrogen. This lesson compares how the Period 3 elements react with water, oxygen and chlorine, and links each pattern to the trends of Lesson 1.

Definition

Reducing agent: a substance that gives electrons to another and is itself oxidised. Example: sodium, $\mathrm{Na \rightarrow Na^+ + e^-}$ , reduces water to hydrogen.

Oxidising agent: a substance that accepts electrons and is itself reduced. Example: chlorine, $\mathrm{Cl_2 + 2e^- \rightarrow 2Cl^-}$ .

Protective oxide layer (passivation): a thin, tightly bound film of oxide on a metal surface that stops further reaction. Example: the layer of $\mathrm{Al_2O_3}$ , a few nanometres thick, on aluminium.

Reactions with water

Element	Conditions and observations	Equation
Na	cold water; violent: melts into a ball, fizzes, moves about; colourless solution, pH 13–14	$\mathrm{2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)}$
Mg	cold water: very slow, a few bubbles; weakly alkaline (pH about 9–10)	$\mathrm{Mg(s) + 2H_2O(l) \rightarrow Mg(OH)_2(s) + H_2(g)}$
Mg	steam: burns with a bright white glow, white solid	$\mathrm{Mg(s) + H_2O(g) \rightarrow MgO(s) + H_2(g)}$
Al	no visible reaction: the oxide layer protects it; powdered aluminium reacts slowly with steam when strongly heated	$\mathrm{2Al(s) + 3H_2O(g) \rightarrow Al_2O_3(s) + 3H_2(g)}$
Si, P, S	no reaction with water under ordinary conditions	none
$\mathrm{Cl_2}$	dissolves and partly reacts: pale green, acidic solution (pH about 2–3) that bleaches litmus	$\mathrm{Cl_2(g) + H_2O(l) \rightleftharpoons HCl(aq) + HClO(aq)}$
Ar	no reaction	none

The metals act as reducing agents, and their reactivity falls from Na to Al because more energy is needed to remove more electrons ( $\mathrm{Na^+}$ : one; $\mathrm{Mg^{2+}}$ : two; $\mathrm{Al^{3+}}$ : three). Chlorine, at the other end, behaves as an oxidising agent; in water it is both oxidised (to +1 in HClO) and reduced (to −1 in HCl).

Reactions with oxygen

Element	Observations on heating in oxygen	Equation
Na	burns with a yellow-orange flame; white solid	$\mathrm{4Na(s) + O_2(g) \rightarrow 2Na_2O(s)}$ (some $\mathrm{Na_2O_2}$ in excess oxygen)
Mg	dazzling white flame; white powder	$\mathrm{2Mg(s) + O_2(g) \rightarrow 2MgO(s)}$
Al	powder burns with white sparks; white solid	$\mathrm{4Al(s) + 3O_2(g) \rightarrow 2Al_2O_3(s)}$
Si	reacts only when strongly heated; white solid	$\mathrm{Si(s) + O_2(g) \rightarrow SiO_2(s)}$
P	white phosphorus catches fire in air at about 30 °C; white flame, dense white smoke	$\mathrm{P_4(s) + 5O_2(g) \rightarrow P_4O_{10}(s)}$ ( $\mathrm{P_4O_6}$ if oxygen is limited)
S	melts and burns with a pale blue flame; choking colourless gas	$\mathrm{S(s) + O_2(g) \rightarrow SO_2(g)}$ ; $\mathrm{SO_3}$ needs a catalyst
Cl, Ar	no direct reaction (chlorine oxides are made indirectly)	none

Sulfur burning on a stone. Top: in daylight, the sulfur has melted to a blood-red liquid. Bottom: in the dark, the pale blue flame of the burning sulfur shows clearly; the product is sulfur dioxide gas. Photo: Johannes 'volty' Hemmerlein, CC BY-SA 3.0, via Wikimedia Commons

Reactions with chlorine

All the elements from sodium to sulfur combine directly with chlorine when heated in a stream of the dry gas:

\mathrm{2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)} \qquad \mathrm{Mg(s) + Cl_2(g) \rightarrow MgCl_2(s)}

\mathrm{2Al(s) + 3Cl_2(g) \rightarrow 2AlCl_3(s)} \qquad \mathrm{Si(s) + 2Cl_2(g) \rightarrow SiCl_4(l)}

\mathrm{P_4(s) + 6Cl_2(g) \rightarrow 4PCl_3(l)}\ \ \text{(limited Cl}_2\text{)} \qquad \mathrm{P_4(s) + 10Cl_2(g) \rightarrow 4PCl_5(s)}\ \ \text{(excess Cl}_2\text{)}

\mathrm{2S(s) + Cl_2(g) \rightarrow S_2Cl_2(l)}

Sodium burns in chlorine with a bright yellow flame, giving white sodium chloride. The products change from white ionic solids (NaCl, $\mathrm{MgCl_2}$ ) to volatile covalent liquids ( $\mathrm{SiCl_4}$ , $\mathrm{PCl_3}$ , $\mathrm{S_2Cl_2}$ ): Lesson 4 explains why.

Oxidation states climb across the period

In the highest oxides and chlorides the oxidation state of the Period 3 element equals its number of outer electrons, because all of them are used in bonding to the more electronegative oxygen or chlorine.

Figure 14.4: Oxidation state of the element in the highest oxide of each Period 3 element. Orange: ionic oxides of metals; blue: covalent oxides.

Worked example 1: hydrogen from sodium

0.230 g of sodium reacts completely with water. Calculate the volume of hydrogen formed at 25 °C and 1 atm ( $V_m = 24.47\ \mathrm{dm^3\,mol^{-1}}$ ; Na = 22.99).

Solution. $n(\mathrm{Na}) = \frac{0.230}{22.99} = 1.000 \times 10^{-2}\ \mathrm{mol}$ From $\mathrm{2Na + 2H_2O \rightarrow 2NaOH + H_2}$ , $n(\mathrm{H_2}) = \frac{1}{2} \times 1.000 \times 10^{-2} = 5.00 \times 10^{-3}\ \mathrm{mol}$ . $V = 5.00 \times 10^{-3} \times 24.47 = 0.122\ \mathrm{dm^3} = 122\ \mathrm{cm^3}$

Worked example 2: which phosphorus oxide?

1.24 g of white phosphorus is burned in a closed flask containing 1.60 g of oxygen. Decide which oxide forms, and its mass (P = 30.97, O = 16.00).

Solution. $n(\mathrm{P_4}) = \frac{1.24}{123.88} = 0.0100\ \mathrm{mol} \qquad n(\mathrm{O_2}) = \frac{1.60}{32.00} = 0.0500\ \mathrm{mol}$ The mole ratio $\mathrm{O_2 : P_4} = 5 : 1$ , exactly that of $\mathrm{P_4 + 5O_2 \rightarrow P_4O_{10}}$ . (For $\mathrm{P_4O_6}$ only 3 mol of $\mathrm{O_2}$ per mol of $\mathrm{P_4}$ would be used.) So $\mathrm{P_4O_{10}}$ forms: $m = 0.0100 \times 283.88 = 2.84\ \mathrm{g}$ which equals $1.24 + 1.60$ , as conservation of mass requires.

Warning

Common mistakes. (1) Writing $\mathrm{MgO}$ as the product of magnesium with cold water: cold water gives $\mathrm{Mg(OH)_2}$ ; steam gives $\mathrm{MgO}$ . (2) Concluding that aluminium is unreactive: it is a reactive metal hidden under an oxide layer. (3) Writing $\mathrm{SO_3}$ as the product of burning sulfur: burning gives $\mathrm{SO_2}$ . (4) Writing phosphorus as P in equations for the element: use $\mathrm{P_4}$ .

Safety. The sodium and water demonstration is done by the teacher, with a piece no bigger than a rice grain, behind a safety screen, with goggles on. Do not look straight at burning magnesium: its light contains ultraviolet. Burning sulfur gives toxic sulfur dioxide and chlorine is toxic: both are handled only in a fume cupboard. If chlorine or sulfur dioxide is breathed in, move the person to fresh air and get medical help.

Tip

Left gives, right takes. Na, Mg and Al give electrons (reducing agents; reactivity Na > Mg > Al). Cl takes electrons (oxidising agent). In the highest oxides, the oxidation state counts the outer electrons: +1, +2, +3, +4, +5, +6, +7.

Bale's Chemistry: Reactions of Period 3 Elements (6 min). Watch for sodium and magnesium with water and steam, and the formulas of the oxides formed when sodium to sulfur react with oxygen.

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Practice

Write balanced equations, with state symbols, for the reaction of (a) sodium with water and (b) magnesium with steam.
Describe what you would see when magnesium burns in oxygen and when sulfur burns in oxygen.
Explain why aluminium does not appear to react with water.
Write equations for the reaction of chlorine with (a) aluminium and (b) silicon.
Give the oxidation states of chlorine in $\mathrm{Cl_2}$ , HCl and HClO, and explain why the reaction of chlorine with water is called a disproportionation.
0.486 g of magnesium reacts completely with steam. Calculate the mass of magnesium oxide formed and the volume of hydrogen at 25 °C and 1 atm (Mg = 24.31, O = 16.00; $V_m = 24.47\ \mathrm{dm^3\,mol^{-1}}$ ).
2.70 g of aluminium burns completely in oxygen. Calculate the mass of oxide formed (Al = 26.98).
Sodium reacts violently with cold water, magnesium very slowly and aluminium hardly at all. Explain this trend in terms of ionisation energy.

Show answersHide answers

(a) $\mathrm{2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)}$ ; (b) $\mathrm{Mg(s) + H_2O(g) \rightarrow MgO(s) + H_2(g)}$ .
Magnesium: a dazzling white flame, leaving a white powder (MgO). Sulfur: melts and burns with a pale blue flame, giving a colourless, choking gas ( $\mathrm{SO_2}$ ).
A thin, unbroken layer of $\mathrm{Al_2O_3}$ on the surface stops water reaching the metal.
(a) $\mathrm{2Al + 3Cl_2 \rightarrow 2AlCl_3}$ ; (b) $\mathrm{Si + 2Cl_2 \rightarrow SiCl_4}$ .
0, −1 and +1. The same element is both reduced (0 to −1) and oxidised (0 to +1) in one reaction.
$n(\mathrm{Mg}) = \frac{0.486}{24.31} = 0.0200\ \mathrm{mol}$ ; MgO: $0.0200 \times 40.31 = 0.806\ \mathrm{g}$ ; $\mathrm{H_2}$ : $0.0200 \times 24.47 = 0.489\ \mathrm{dm^3}$ .
$n(\mathrm{Al}) = \frac{2.70}{26.98} = 0.1001\ \mathrm{mol}$ ; $n(\mathrm{Al_2O_3}) = 0.0500\ \mathrm{mol}$ ; mass $= 0.0500 \times 101.96 = 5.10\ \mathrm{g}$ .
To react, each atom must lose its outer electrons. Na loses one (IE 496 kJ mol⁻¹); Mg must lose two (738 + 1451 kJ mol⁻¹) and Al three (578 + 1817 + 2745 kJ mol⁻¹). The energy cost rises, so the reaction becomes harder; aluminium is also protected by its oxide layer.

Quick check

What forms when magnesium reacts with steam?

Key points

With water: Na violently (NaOH + $\mathrm{H_2}$ ); Mg slowly with cold water, quickly with steam (MgO + $\mathrm{H_2}$ ); Al protected by its oxide; Si, P, S not at all; $\mathrm{Cl_2}$ disproportionates to HCl + HClO.
With oxygen: $\mathrm{Na_2O}$ , MgO, $\mathrm{Al_2O_3}$ , $\mathrm{SiO_2}$ , $\mathrm{P_4O_{10}}$ (or $\mathrm{P_4O_6}$ ), $\mathrm{SO_2}$ .
With chlorine: NaCl, $\mathrm{MgCl_2}$ , $\mathrm{AlCl_3}$ , $\mathrm{SiCl_4}$ , $\mathrm{PCl_3}$ or $\mathrm{PCl_5}$ , $\mathrm{S_2Cl_2}$ .
The metals are reducing agents (reactivity Na > Mg > Al); chlorine is an oxidising agent. The oxidation state in the highest oxide rises from +1 (Na) to +7 (Cl).

Sources and credits

Image: Burning-sulfur.png by Johannes 'volty' Hemmerlein, via Wikimedia Commons, licensed under CC BY-SA 3.0.
Image: Phosphorus Burning in Oxygen.JPG by Leiem, via Wikimedia Commons, licensed under CC BY-SA 4.0.
Ionisation energies: Molar ionization energies of the elements, Wikipedia contributors, licensed under CC BY-SA 4.0.
Video: Reactions of Period 3 elements, Paper 1 - AQA A Level Chemistry, Bale's Chemistry, embedded from YouTube.
The text, tables, chart, worked examples and Practice items were written and drawn by Ganzaa.
This lesson uses CC BY-SA material and is shared under CC BY-SA 4.0.

1Same water, three different answers

2Reactions with water

3Reactions with oxygen

4Reactions with chlorine

5Oxidation states climb across the period

Same water, three different answers

Reactions with water

Reactions with oxygen

Reactions with chlorine

Oxidation states climb across the period