Trends in Chemical Properties of Group 15 Elements and Their Compounds

Lessons

Senior 4 · Chemistry

Trends in Chemical Properties of Group 15 Elements and Their Compounds

5 lessons

Lesson 2 of 5

Reactions of the Group 15 Elements, Their Hydrides and Oxides

25 min

Fertiliser from a thunderstorm

A flash of lightning heats the air it passes through to many thousands of degrees, hot enough to force nitrogen and oxygen to combine. The nitrogen monoxide formed is oxidised further and washed down by rain as nitric acid and nitrates, a small natural supply of fixed nitrogen for the soil. The same reaction happens, unwanted, inside car engines, and the nitrogen oxides from exhausts help to form smog and acid rain. Nitrogen reacts only under extreme conditions; phosphorus, arsenic, antimony and bismuth react much more readily. This lesson surveys their reactions and the properties of their hydrides and oxides.

Definition

Disproportionation: a reaction in which the same element is both oxidised and reduced. Example: $\mathrm{P_4 + 3NaOH + 3H_2O \rightarrow PH_3 + 3NaH_2PO_2}$ : phosphorus (0) goes to −3 in phosphine and to +1 in sodium phosphinate.

Hydride (of a non-metal): a binary compound of the element with hydrogen. Example: ammonia, $\mathrm{NH_3}$ ; phosphine, $\mathrm{PH_3}$ .

Reactions of the elements

With oxygen. Nitrogen combines with oxygen only at very high temperature (lightning, engines, furnaces), and the reaction is endothermic; the NO formed reacts with more oxygen in cool air to give brown nitrogen dioxide:

\mathrm{N_2(g) + O_2(g) \rightleftharpoons 2NO(g)} \quad \Delta H = +180.6\ \mathrm{kJ\,mol^{-1}} \qquad \mathrm{2NO(g) + O_2(g) \rightarrow 2NO_2(g)}

Phosphorus burns brightly. With limited oxygen it gives phosphorus(III) oxide, with excess oxygen phosphorus(V) oxide, both made of molecules based on the $\mathrm{P_4}$ tetrahedron:

\mathrm{P_4(s) + 3O_2(g) \rightarrow P_4O_6(s)} \qquad \mathrm{P_4(s) + 5O_2(g) \rightarrow P_4O_{10}(s)}

Arsenic and antimony burn to the +3 oxides $\mathrm{As_4O_6}$ and $\mathrm{Sb_4O_6}$ ; bismuth gives $\mathrm{Bi_2O_3}$ . The +3 state becomes more favoured down the group (the inert pair effect again).

With metals. Heated reactive metals form nitrides and phosphides, which are hydrolysed by water to the hydrides:

\mathrm{3Mg + N_2 \rightarrow Mg_3N_2} \qquad \mathrm{Mg_3N_2 + 6H_2O \rightarrow 3Mg(OH)_2 + 2NH_3}

\mathrm{Ca_3P_2 + 6H_2O \rightarrow 3Ca(OH)_2 + 2PH_3}

With chlorine. Nitrogen does not react directly. Phosphorus burns in chlorine to give phosphorus trichloride, and in excess chlorine phosphorus pentachloride:

\mathrm{P_4 + 6Cl_2 \rightarrow 4PCl_3} \qquad \mathrm{PCl_3 + Cl_2 \rightleftharpoons PCl_5}

Arsenic, antimony and bismuth give trichlorides, $\mathrm{MCl_3}$ .

With alkali. White phosphorus disproportionates when warmed with sodium hydroxide solution, giving phosphine, a poisonous gas that may catch fire in air: $\mathrm{P_4 + 3NaOH + 3H_2O \rightarrow PH_3 + 3NaH_2PO_2}$ .

With acids. Bismuth, the most metallic element, dissolves in nitric acid to give a salt, bismuth(III) nitrate; the non-metals are not attacked by dilute acids, though hot concentrated nitric acid oxidises phosphorus to phosphoric acid.

The hydrides

Table 10.3: Hydrides of the Group 15 elements

Hydride	Boiling point / °C	As a base	Thermal stability
ammonia, $\mathrm{NH_3}$	−33.3	weak base; forms $\mathrm{NH_4^+}$ salts readily	stable
phosphine, $\mathrm{PH_3}$	−87.7	very weak base	decomposes on heating
arsine, $\mathrm{AsH_3}$	−62.5	almost no basic character	decomposes on gentle heating
stibine, $\mathrm{SbH_3}$	−17	none	decomposes slowly at room temperature
bismuthine, $\mathrm{BiH_3}$	16.8	none	very unstable

Boiling points rise from $\mathrm{PH_3}$ to $\mathrm{BiH_3}$ , because bigger molecules with more electrons have stronger London forces. Ammonia breaks the trend: it boils about 55 °C higher than phosphine, because nitrogen is small and very electronegative and has a lone pair, so ammonia molecules form hydrogen bonds to each other (Figure 10.4).
Basic strength falls down the group. The lone pair on the small nitrogen atom is concentrated and readily accepts a proton; on the larger atoms it is spread over a bigger volume and is less available.
Thermal stability falls as the E–H bond gets longer and weaker; the hydrides become stronger reducing agents down the group.

Figure 10.4: Boiling points of the Group 15 hydrides. Ammonia's value is far above the line through the other four.

The oxides: acid to base

Table 10.4: Acid–base character of the oxides

Element	Oxides	Character
Nitrogen	$\mathrm{N_2O_3}$ , $\mathrm{N_2O_5}$ (and neutral $\mathrm{N_2O}$ , NO)	acidic: give nitrous and nitric acids
Phosphorus	$\mathrm{P_4O_6}$ , $\mathrm{P_4O_{10}}$	acidic: give phosphonic and phosphoric acids
Arsenic	$\mathrm{As_4O_6}$	amphoteric, mainly acidic
Antimony	$\mathrm{Sb_4O_6}$	amphoteric
Bismuth	$\mathrm{Bi_2O_3}$	basic: dissolves in acids to give $\mathrm{Bi^{3+}}$ salts

The change from acidic to basic follows the change from non-metal to metal, exactly as in Groups 13 and 14. For one element, the oxide with the higher oxidation state is the more acidic: $\mathrm{P_4O_{10}}$ is more acidic than $\mathrm{P_4O_6}$ .

\mathrm{P_4O_{10}(s) + 6H_2O(l) \rightarrow 4H_3PO_4(aq)} \qquad \mathrm{Bi_2O_3(s) + 6HNO_3(aq) \rightarrow 2Bi(NO_3)_3(aq) + 3H_2O(l)}

Worked example 1: ammonia from a nitride

2.02 g of magnesium nitride is added to excess water. Calculate the volume of ammonia formed at room temperature and pressure (24.0 dm³ mol⁻¹). $A_{r}$ : Mg 24.31, N 14.01.

Solution. $\mathrm{Mg_3N_2 + 6H_2O \rightarrow 3Mg(OH)_2 + 2NH_3}$ . $M(\mathrm{Mg_3N_2}) = 3(24.31) + 2(14.01) = 100.95\ \mathrm{g\,mol^{-1}} \qquad n = \frac{2.02}{100.95} = 0.0200\ \mathrm{mol}$ $n(\mathrm{NH_3}) = 2 \times 0.0200 = 0.0400\ \mathrm{mol} \qquad V = 0.0400 \times 24.0 = 0.960\ \mathrm{dm^3}$ The gas turns damp red litmus paper blue, which shows that the nitride contained $\mathrm{N^{3-}}$ .

Worked example 2: burning phosphorus

6.20 g of white phosphorus burns completely in excess oxygen. (a) Write the equation and calculate the mass of oxide formed. (b) The oxide is dissolved in water. Name the acid formed and give its formula. $A_{r}$ : P 30.97, O 16.00.

Solution. (a) $\mathrm{P_4 + 5O_2 \rightarrow P_4O_{10}}$ . $M(\mathrm{P_4}) = 123.88 \qquad n = \frac{6.20}{123.88} = 0.0500\ \mathrm{mol}$ $M(\mathrm{P_4O_{10}}) = 4(30.97) + 10(16.00) = 283.88 \qquad m = 0.0500 \times 283.88 = 14.2\ \mathrm{g}$ (b) Phosphoric(V) acid, $\mathrm{H_3PO_4}$ : $\mathrm{P_4O_{10} + 6H_2O \rightarrow 4H_3PO_4}$ . An acidic oxide gives an acid.

Warning

Common mistakes. (1) Writing " $\mathrm{N_2 + 2O_2 \rightarrow 2NO_2}$ " as a direct reaction: nitrogen first forms NO, which then reacts with more oxygen. (2) Writing $\mathrm{P_2O_5}$ in equations without saying it is the empirical formula of $\mathrm{P_4O_{10}}$ . (3) Explaining ammonia's high boiling point by its "larger mass": $\mathrm{NH_3}$ is the lightest hydride; the cause is hydrogen bonding. (4) Calling phosphine a strong base because ammonia is a base: basicity falls sharply down the group.

Safety. Phosphine, arsine and stibine are extremely toxic gases; they are never made in school laboratories. Nitrogen dioxide is a toxic, brown gas: reactions that produce it are done in a fume cupboard. If a chemical gets in your eyes, rinse with plenty of running water and get medical help.

Tip

Down Group 15 everything follows "non-metal to metal": oxides acidic → basic; hydrides less basic and less stable; +3 more stable than +5. The one big exception is ammonia's hydrogen bonding, which gives it a high boiling point and makes it extremely soluble in water.

Periodic Videos: More White Phosphorus (6 min). Watch for how white phosphorus is handled and for phosphorus pentoxide, the white oxide (P₄O₁₀) that forms when it burns.

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Practice

Write equations for: (a) nitrogen reacting with oxygen in lightning; (b) phosphorus burning in a limited supply of oxygen; (c) phosphorus trichloride reacting with more chlorine.
Explain why the reaction of nitrogen with oxygen needs a very high temperature.
Explain why ammonia has a higher boiling point than phosphine, although phosphine has the greater molar mass.
Predict the boiling point trend from phosphine to bismuthine, and explain it.
Arrange $\mathrm{P_4O_{10}}$ , $\mathrm{Sb_4O_6}$ and $\mathrm{Bi_2O_3}$ in order of increasing basic character. Write an equation showing that $\mathrm{Bi_2O_3}$ is basic.
Identify the oxidation numbers of phosphorus in the reaction $\mathrm{P_4 + 3NaOH + 3H_2O \rightarrow PH_3 + 3NaH_2PO_2}$ , and explain why it is called disproportionation.
Calculate the volume of phosphine, at room temperature and pressure, formed when 9.10 g of calcium phosphide, $\mathrm{Ca_3P_2}$ , reacts with excess water. $A_{r}$ : Ca 40.08, P 30.97; 24.0 dm³ mol⁻¹.

Show answersHide answers

(a) $\mathrm{N_2 + O_2 \rightarrow 2NO}$ ; (b) $\mathrm{P_4 + 3O_2 \rightarrow P_4O_6}$ ; (c) $\mathrm{PCl_3 + Cl_2 \rightleftharpoons PCl_5}$ .
The N≡N bond (945 kJ mol⁻¹) gives a very high activation energy, and the reaction is endothermic (+180.6 kJ mol⁻¹), so it only happens, and only gives a useful yield, at very high temperatures.
Ammonia molecules form hydrogen bonds (small, highly electronegative N with a lone pair and N–H bonds); phosphine has only weak London and dipole forces.
It rises (−87.7, −62.5, −17, 16.8 °C): the molecules have more electrons, so London forces are stronger.
$\mathrm{P_4O_{10}}$ < $\mathrm{Sb_4O_6}$ < $\mathrm{Bi_2O_3}$ . $\mathrm{Bi_2O_3 + 6HNO_3 \rightarrow 2Bi(NO_3)_3 + 3H_2O}$ .
0 in $\mathrm{P_4}$ ; −3 in $\mathrm{PH_3}$ ; +1 in $\mathrm{NaH_2PO_2}$ . The same element is both reduced (0 to −3) and oxidised (0 to +1).
$M(\mathrm{Ca_3P_2}) = 182.18$ ; $n = \frac{9.10}{182.18} = 0.0500$ mol; $\mathrm{Ca_3P_2 + 6H_2O \rightarrow 3Ca(OH)_2 + 2PH_3}$ , so $n(\mathrm{PH_3}) = 0.100$ mol; $V = 0.100 \times 24.0 = 2.40\ \mathrm{dm^3}$ .

Quick check

Ammonia boils at −33.3 °C but phosphine, a heavier molecule, boils at −87.7 °C. Why?

Key points

Nitrogen combines with oxygen only at very high temperature ( $\mathrm{N_2 + O_2 \rightleftharpoons 2NO}$ , +180.6 kJ mol⁻¹); phosphorus burns readily to $\mathrm{P_4O_6}$ or $\mathrm{P_4O_{10}}$ ; As, Sb and Bi give +3 oxides.
Reactive metals form nitrides and phosphides that water hydrolyses to $\mathrm{NH_3}$ and $\mathrm{PH_3}$ ; phosphorus forms $\mathrm{PCl_3}$ and $\mathrm{PCl_5}$ ; white phosphorus disproportionates in alkali.
Hydrides: basic strength and thermal stability fall from $\mathrm{NH_3}$ to $\mathrm{BiH_3}$ ; boiling points rise from $\mathrm{PH_3}$ to $\mathrm{BiH_3}$ , with $\mathrm{NH_3}$ anomalously high because of hydrogen bonding.
Oxides change from acidic (N, P) through amphoteric (As, Sb) to basic (Bi).

Sources and credits

Image: Lightning over Oradea Romania 3.jpg by Mircea Madau, via Wikimedia Commons, public domain.
Image: White phosphorus.jpg by BXXXD, via Wikimedia Commons, licensed under CC BY-SA 4.0.
Data: boiling points from the Wikipedia articles Ammonia, Phosphine, Arsine, Stibine and Bismuthine; enthalpy of formation of NO from Nitric oxide; all CC BY-SA 4.0, adapted by Ganzaa.
First aid for chemical splashes: NIOSH Pocket Guide, first aid, CDC.
Video: More White Phosphorus - Periodic Table of Videos, Periodic Videos (University of Nottingham), embedded from YouTube.
The text, chart, tables, worked examples and Practice items were written by Ganzaa.
This lesson adapts CC BY-SA material and is shared under CC BY-SA 4.0.

1Fertiliser from a thunderstorm

2Reactions of the elements

3The hydrides

4The oxides: acid to base

Fertiliser from a thunderstorm

Reactions of the elements

The hydrides

The oxides: acid to base