1. Group 5 Members General comment Elements symbol Nitrogen N
Phosphorus P
General comment
Thus nitrogen and phosphorus are non metals
Nitrogen forms stable NN bonds that make nitrogen inert.
Nitrogen and phosphorus form compounds with valency 3 using the three unpaired p-electrons

2. Group 5 General comment Trial 1
Phosphorus is able forms compound with valency 3 or 5
To form compound of valency 5 it promotes one of the 3s-electrons to a 3d-orbital.
i.e. 3s23p33d s13p33d1
However, Nitrogen does not form compound with a valence of 5 because it lacks a d-orbital.
Trial 1
(a) Write the full electronic configurations of nitrogen and phosphorus atoms.(2 marks)
(b) Explain why phosphorus (V) chloride is formed but NOT nitrogen (V) chloride.
(1 mark)

3. Group 5 Nitrogen Nitrogen Laboratory preparation
Industrial preparation
By factional distillation of air from which carbon dioxide and water vapour has been removed.
Nitrogen
Laboratory preparation
By passing of ammonia over hot copper oxide
2NH3 (g) + 3CuO(s)
→ N2(g) + 3H2O(l) + 3Cu(s)

4. Group 5 Nitrogen Ammonia Use as a cooling agent.
used to make fertilizer
to make nitric acid
Compound of nitrogen
Ammonia
Industrial preparation
It is prepared in the Haber process by reacting hydrogen with nitrogen
N2(g) + 3H2(g)
2NH3(g) ΔH = -ve

5. Group 5 Favorable condition for production of ammonia
The reaction of nitrogen and hydrogen to form ammonia is exothermic and results in decrease in gas volume.
Le Chatelier’s principle predicts, therefore, that the yield of ammonia will be greatest at low temperature and high pressure
Favorable condition for production of ammonia
because the reaction proceeds with a decrease in the number of moles of a gas
At low temperature, however, equilibrium is reached slowly; at high pressure, the cost of equipment is high.

6. Group 5 (a) group 1 elements It forms amides and hydrogen
Favorable condition for production of ammonia
In practice a compromise has to be struck. Temperature of C is used at pressure of 1 x 102 – 1 x 103 atmospheres and the rate of the reaction increased by use of an iron catalyst.
Reactions of ammonia
(a) group 1 elements
It forms amides and hydrogen
2Na(s) + 2NH3(g)
→ 2NaNH2(s) + H2(g)

7. Group 5 (a) group 1 elements It forms amides and hydrogen
Favorable condition for production of ammonia
In practice a compromise has to be struck. Temperature of C is used at pressure of 1 x 102 – 1 x 103 atmospheres and the rate of the reaction increased by use of an iron catalyst.
Reactions of ammonia
(a) group 1 elements
It forms amides and hydrogen
2Na(s) + 2NH3(g)
→ 2NaNH2(s) + H2(g)

8. Group 5 Reactions of ammonia (b) As a reducing agent
(i) it reduces copper and lead oxides to metals
2NH3 (g) + 3CuO(s)
→ N2 (g) + 3H2O (l) + 3Cu(s)
2NH3 (g) + 3PbO(s)
→ N2 (g) + 3H2O (l) + 3Pb(s)
Reactions of ammonia
(b) As a reducing agent
(ii) It is oxidized by chlorine and sodium chlorate (I) at room temperature to nitrogen.
2NH3 (g) + 3Cl2 (g)
→ N2(g) + 6HCl(g)
(NH4Cl is formed with excess NH3)

9. Group 5 Reactions of ammonia (c) Reaction with oxygen Trial 2
Ammonia burns in air with a pale green flame to give nitrogen and water.
4NH3 (g) + 3O2 (g)
→ 2N2(g) + 6H2O(l)
Trial 2
Compare the Chemistry of ammonia with that of water by describing and interpreting their reactions with
(i) sodium
(ii) calcium chloride solution
(iii) chlorine
(iv) anhydrous copper (II) sulphate solution.

10. Group 5 Ammonia salts and their uses (a) Ammonium sulphate (NH4)2 SO4
It is made from ammonia and sulphuric acid or by passing ammonia and carbon dioxide into a suspension of calcium sulphate in water and filtering off the calcium carbonate.
It is used as a fertilizer and in more concentrated form as a weed killer.
Ammonia salts and their uses
(b) Ammonium chloride, NH4Cl ( Sal- ammoniac).
Is made by reaction between ammonia and hydrochloric acid, or by crystallization from a mixed solution of ammonium sulphate and sodium chloride. Alternatively, ammonium sulphate and sodium chloride are heated in the solid state and ammonium chloride collects as a sublimate.

11. Group 5 Ammonia salts and their uses Ammonia salts and their uses
(b) Ammonium chloride, NH4Cl ( Sal- ammoniac).
It is used in dry cells and as a flux soldering, tin plating and galvanizing.
Ammonia salts and their uses
c) Ammonium carbonate,
(NH4)2CO3
Commercial ammonium carbonate is made by heating a mixture of solid calcium carbonate and ammonium sulphate. The white solid collects as a sublimate is a mixture of ammonium hydrogen carbonate, ammonium carbamate and ammonium carbonate.

12. Group 5 Ammonia salts and their uses Ammonia salts and their uses
c) Ammonium carbonate, (NH4)2CO3
Commercial ammonium carbonate is made by heating a mixture of solid calcium carbonate and ammonium sulphate. The white solid collects as a sublimate is a mixture of ammonium hydrogen carbonate, ammonium carbamate and ammonium carbonate.
Ammonia salts and their uses
(d) Ammonium nitrate, NH4NO3
It is made by direct combination between ammonia and nitric acid. It decomposes on heating, explosively if the temperature gets too high.
NH4NO3 (s) → N2O (g) + 2H2O (g)
 It is used as a fertilizer and in making explosives:

13. Group 5 Ammonia salts and their uses Ammonia salts and their uses
Ammonium molybdate (IV), (NH4)2MoO4.
This is made by dissolving molybdenum (VI) oxide in ammonia solution and treating a solution of the crystals obtained with ammonia. It is used as a test for phosphates.
Ammonia salts and their uses
(d) Ammonium nitrate, NH4NO3
It is made by direct combination between ammonia and nitric acid. It decomposes on heating, explosively if the temperature gets too high.
NH4NO3 (s) → N2O (g) + 2H2O (g)
 It is used as a fertilizer and in making explosives:

14. Group 5 Trial 3 Nitric acid
Stating the appropriate conditions, describe what happens when ammonia reacts with
Oxygen
Sodium
Chlorine
Silver nitrate
Carbon dioxide
Nitric acid
Industrial preparation
is made by catalytic oxidation of ammonia gas with excess air in the presence of a platinum- rhodium catalyst.
4NH3 (g) + 5O2(g)
→ 4NO (g) + 6H2O (g)
Then
2NO(g) + O2(g) → 2NO2(g)

15. Group 5 Nitric acid Industrial preparation
Then nitrogen dioxide is dissolved in water in presence oxygen to form nitric acid
4NO2 (g) + 2H2O (l) + O2(g)
→ 4HNO3 (aq)
Nitric acid
Uses
i. synthesis dyes
ii. Synthesis of fertilizers.
iii. Synthesis of drugs
iv. Organic synthesis

16. Group 5 Reactions of Nitric acid (a) As an acid
turns blue litmus paper red
Liberates carbon dioxide from carbonate
Reacts with metals to produce salts and hydrogen
Mg(s) + 2HNO3(aq)
→ Mg(NO3)2(aq) + H2(g)
Reactions of Nitric acid
(b) as oxidizing agent
it oxidizes carbon to carbon dioxide.
C(s) + 4HNO3(aq)
→ CO2(g) + 4NO2(g) + 2H2O(l)

17. Group 5 Reactions of Nitric acid (b) as oxidizing agent
(ii) It oxidizes Sulphur and phosphorus to sulphuric and phosphoric acids respectively
S (s) + 6HNO3 (aq)
→ H2SO4(aq) + 6NO2(g) + 2H2O(l)
P (s) + 5HNO3 (aq)
→ H3PO4(aq) + 5NO2(g) + H2O(l)
Reactions of Nitric acid
(b) as oxidizing agent
(iii) 50% nitric acid oxidizes copper to copper nitrate and nitrogen monoxide.
3Cu (s) + 8HNO3 (aq) →
3Cu(NO3)2(aq) + 4H2O(l) + 2NO(g)

18. Group 5 Reactions of Nitric acid (b) as oxidizing agent
(iv) Concentrated nitric acid oxidizes copper to copper nitrate and nitrogen dioxide.
Cu (s) + 4HNO3 (l) →
Cu(NO3)2(aq) + 2H2O (l) + 2NO2(g)
Trial 4
Describe how nitric acid:
(i) is manufactured. (9 marks)
(ii) reacts with copper (your answer should include equations for the reaction) (5 marks)
(iii) State two industrial uses of nitric acid. (2 marks)

19. Group 5 Reactions of Nitric acid (b) as oxidizing agent
(v) Oxidizes acidified iron II salts to iron III salts.
3Fe2+(aq) + 4H+(aq) + NO3-(aq) →
3Fe3+(aq) + 2H2O(l) + NO(g)
Reactions of Nitric acid
(b) as oxidizing agent
(vi) Nitric acid oxidizes hydrogen sulphide and hydrogen chloride to sulphur and chlorine respectively.
3H2S (l) + 2HNO3 (aq)
→ 4H2O (l) + 2NO (g) + 3S (s)
3HCl (g) + HNO3 (aq)
→ 2H2O(l) + NOCl(g) + Cl2(g)

20. Group 5
Trial 5
State what is observed and write ionic equations for the reactions that take place between nitric acid and
(i) Iron II sulphate
(ii) Hydrogen sulphide
(iii) hydrogen chloride gas
Nitrates
Preparation
Nitrates can be prepared by reaction of metals, oxides or hydroxides with nitric acid.
Mg (s) + 2HNO3 (aq)
→ Mg (NO3)2 (aq) + H2 (g)
Na2O (s) + 2HNO3 (aq)
→ 2NaNO3 (aq) + H2O (l)
NH4OH (aq) + HNO3 (aq)
→ NH4NO3 (aq) + H2O (l)

21. Group 5 Effect of heat on nitrates
(i) Sodium and potassium nitrates decompose on heating into the nitrites and oxygen.
2KNO3 (s) heat 2KNO2(s) + O2(g)
Effect of heat on nitrates
(ii) Most of the nitrates give nitrogen dioxide, oxygen and metallic oxides unless the latter decompose on heating into the metal and oxygen.
2Mg (NO3)2 (s) heat
2MgO(s) + 4NO2(g) + O2(g)
2AgNO3 (s) heat
2Ag (s) + 2NO2(g) + O2(g)

22. Group 5 Effect of heat on nitrates
(iii) Ammonium nitrate gives dinitrogen oxide and water when cautiously heated.
NH4NO3 (s) heat
N2O (g) + 2H2O (l)
Effect of heat on nitrates
Solutions of nitrates are reduced by zinc or, better, by Dewarda's alloy (Al, Cu, Zn) in alkaline solution to ammonia.
NO3-(aq) + 4Zn(s) + 7OH-(aq) + 6H2O (l)
→ NH3(g) + 4[Zn(OH)4]2-(aq)
The reaction can be used quantitatively for estimating nitrates by passing the ammonia formed into a measured excess of standard acid.

23. Group 5 Testing for nitrates Testing for nitrates
Some nitrates decompose on heating to give brown fumes of nitrogen dioxide
Nitrate in acidic media oxide copper to copper nitrates and brown fumes of nitrogen dioxide
Testing for nitrates
iii. Nitrates are reduced by zinc to ammonia that turn moist red litmus paper blue.
iv. The nitrate and the sulphuric acid react to form some nitric acid and some of this is then reduced by the iron (II) to form nitrogen monoxide. The formed nitrogen oxide reacts with the iron (II) ions to form a brown complex [Fe(H2O)5(NO)]2+ as a brown ring.

24. Group 5 Testing for nitrates Brown ring Phosphorus
Electron configuration
1s22s22p63s23p3

25. Group 5 Phosphorus Uses Phosphorus Chemical properties
Phosphorus is used in the match industry, in making rat poison, and in making smoke bombs. About 80% of it is converted to phosphoric acid, H3PO4.
Phosphorus
Chemical properties
Reaction with sodium hydroxide
Phosphorus reacts with hot concentrated sodium hydroxide to form phosphine
P4 (s) + 3NaOH (aq) + 3H2O (l)
→ 3NaH2PO2(aq) + PH3(g)

26. Group 5 Halides of Phosphorus 3-valent excited state, 5
Phosphorus, unlike nitrogen, can form halides of valences 3 and 5. To form compounds of valence 5, one 3s electrons is promoted to a vacant 3d- orbital to give five unpaired electrons.
3-valent excited state, 5
3s23p33d s13p33d1

27. Group 5 The trihalides, PX3
Phosphorus forms all the four covalent trihalides, PF3, PCl3, PBr3, PI3.
Except the trifluoride, all can be made by direct synthesis.
2P(s) + 3Cl2(g) 2PCl3(l)
The trifluoride is a gas, the trichloride and tribromide are liquids and the triiodide is a solid.
The trihalides, PX3
Phosphorus trichloride is readily hydrolysed by water to phosphonic acid, H3PO3, and hydrogen chloride.
PCl3 (l) + 3H2O (l) →
H3PO3 (aq) + 3HCl (aq)

28. Group 5 The trihalides, PX3
Phosphorus trichloride combines with oxygen and reversibly with chlorine to form phosphorus (V) compounds.
2PCl3 (l) + O2 (g) → 2POCl3 (l)
2PCl3 (l) + 2Cl2 (g) PCl5 (s)
The trihalides, PX3
In Organic Chemistry, phosphorus trichloride is used in the preparation of alkyl chloride and acid chloride from alcohols and carboxylic acids respectively.
 3CH3CH2CH2OH(l) + PCl3 (l) →
3CH3CH2CH2Cl(l) + H3PO3(aq)
3CH3CH2COOH(l) + PCl3(l) →
3CH3CH2COCl + H3PO3(aq)

29. Group 5 Phosphorus pentahalides, PX5
Except the pentaiodide, all the pentahalides can be prepared. Phosphorus pentafluoride is a gas. The pentachloride and pentabromide are solids.
Phosphorus pentahalides, PX5
Phosphorus pentachloride is prepared by passing chlorine through a cold flask into which phosphorus trichloride is dripping.
PCl3 (l) + Cl2 (g) PCl5 (s)

30. Group 5 Oxides of phosphorus, P4O6 and P4O10
Phosphorus pentahalides, PX5
Like the trichloride, it is attacked by compounds containing the hydroxyl groups, e.g,
 PCl5 (s) + H2O (l)
→ POCl3 (aq) + 2HCl (aq)
POCl3 (aq) + 3H2O (l)
→ H3PO4 (aq) + 3HCl (aq)
CH3COOH (l) + PCl5 s)
→ CH3COCl (l) + POCl3 (l) + HCl (g)
Oxides of phosphorus, P4O6 and P4O10
Phosphorus (III) oxide, P4O6
It is a covalent colourless salt (m.pt C and b.pt. 1730C) which is a dimer of phosphorus (III) oxide, P2O3, and it is obtained by burning phosphorus in a limited amount of air:
P4 (s) + 3O2 (g) → P4O6 (s)

31. Group 5 Oxides of phosphorus, P4O6 and P4O10
Phosphorus (V) oxide, P4O10
Phosphorus (V) oxide is formed as a white covalent solid when phosphorus is burnt in excess air.
P4 (s) + 5O2 (g) → P4O10 (s)
Phosphorus (V) oxide, P4O10
It is a strong dehydrating agent
a) It dehydrates nitric and sulphuric acids to dinitrogen pentoxide and sulphur (VI) oxide respectively, acid to acid anhydride.
2HNO3(aq) + P2O5(s)
→ N2O5 + 2HPO3 (aq)
H2SO4(aq) + P2O5(s)
→ SO3 + 2HPO3(aq)

32. Group 5 Phosphorus (V) oxide, P4O10
b) it dehydrates acid amide to nitrile.
CH3CONH2(l) + P2O5(s)
→ CH3CN(g) + HPO3(aq)
Phosphorus (V) oxide, P4O10
c) It dehydrates alcohols to alkenes.
CH3CH2OH(l) + P2O5(s)
→ H2C=CH2(g) + 2HPO3(aq)

33. Group 5 Phosphorus (V) oxide, P4O10
b) it dehydrates carboxylic acid to acid anhydride.
2CH3COOH(l) + P2O5(s) →
(CH3CO)2O(l) + 2HPO3(aq
END
THANX