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Liquid nitrogen VA group. Nitrogen and Phosphorous and their compounds. Lecture 14 PhD Halina Falfushynska.

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Presentation on theme: "Liquid nitrogen VA group. Nitrogen and Phosphorous and their compounds. Lecture 14 PhD Halina Falfushynska."— Presentation transcript:

1 Liquid nitrogen VA group. Nitrogen and Phosphorous and their compounds. Lecture 14 PhD Halina Falfushynska

2 Atomic and Physical Properties of Group 15 Elements

3 Electronic Configuration. The valence shell electronic configuration of these elements is ns2np3. The s orbital in these elements is completely filled and p orbitals are half- filled, making their electronic configuration extra stable. Atomic and Ionic Radii. Covalent and ionic (in a particular state) radii increase in size down the group. There is a considerable increase in covalent radius from N to P. However, from As to Bi only a small increase in covalent radius is observed. This is due to the presence of completely filled d and/or f orbitals in heavier members. Ionisation Enthalpy. Ionisation enthalpy decreases down the group due to gradual increase in atomic size. Because of the extra stable half-filled p orbitals electronic configuration and smaller size, the ionisation enthalpy of the group 15 elements is much greater than that of group 14 elements in the corresponding periods. The order of successive ionisation enthalpies, as expected is ΔH1 < ΔH2 < ΔH3 Physical Properties. All the elements of this group are polyatomic. Dinitrogen is a diatomic gas while all others are solids. Metallic character increases down the group. Nitrogen and phosphorus are non-metals, arsenic and antimony metalloids and bismuth is a metal. This is due to decrease in ionisation enthalpy and increase in atomic size. The boiling points, in general, increase from top to bottom in the group but the melting point increases up to arsenic and then decreases up to bismuth. Except nitrogen, all the elements show allotropy.

4 Preparation of dinitrogen In the laboratory, dinitrogen is prepared by treating an aqueous solution of ammonium chloride with sodium nitrite. NH4CI(aq) + NaNO2 (aq) → N2 (g) + 2H2O(l) + NaCl (aq) It can be obtained by the thermal decomposition of ammonium dichromate. (NH4)2Cr2O7 (Heat) → N2 + 4H2O + Cr2O3 Very pure nitrogen can be obtained by the thermal decomposition of sodium or barium azide. Ba(N3)2 → Ba + 3N2 Air (4N 2 + O 2) + C → 4N 2 + CO 2 NH 3 + 3O 2 → 2N 2 + 6H 2 O 2NH 3 + 3Cl 2 → N 2 + 6HCl

5 Reactivity towards hydrogen All the elements of Group 15 form hydrides of the type EH3 where E = N, P, As, Sb or Bi. N2(g) + 3H2(g) (773 k) ==> 2NH3(g); ΔH= –46.1 kJmol–1 Р 4 + 6Н 2 (heat, p) ==> 4РН 3 The stability of hydrides decreases from NH3 to BiH3. the reducing character of the hydrides increases. Ammonia is only a mild reducing agent while BiH3 is the strongest reducing agent amongst all the hydrides. Basicity also decreases in the order NH3 > PH3 > AsH3 > SbH3 > BiH3. Reactivity towards oxygen All these elements form two types of oxides: E2O3 and E2O5. The oxide in the higher oxidation state of the element is more acidic than that of lower oxidation state. Their acidic character decreases down the group. The oxides of the type E2O3 of nitrogen and phosphorus are purely acidic. N 2 (g) + O 2 (g) (heat) ==> 2NO (g) P 4 + 5O 2 (heat) ==> 2P 4 O 10

6 Reactivity towards halogens These elements react to form two series of halides: EX3 and EX5. Nitrogen does not form pentahalide. Pentahalides are more covalent than trihalides. All the trihalides of these elements except those of nitrogen are stable. In case of nitrogen, only NF3 is known to be stable. P 4 + 6Cl 2 ==> 4PCl 3 3PCl 5 + 2P ==> 5PCl 3 3PCl 5 + P 2 O 5 ==> 5POCl 3 Reactivity towards metals All these elements react with metals to form their binary compounds exhibiting –3 oxidation state, such as, Ca3N2 (calcium nitride) Ca3P2 (calcium phosphide), Na3As2 (sodium arsenide), Zn3Sb2 (zinc antimonide) and Mg3Bi2 (magnesium bismuthide). 3Mg + N 2 ==> Mg 3 N 2 Mg 3 N 2 + 6H 2 O ==> 3Mg(OH) 2 + 2NH 3

7 Ammonia Ammonia is present in small quantities in air and soil where it is formed by the decay of nitrogenous organic matter e.g., urea. NH 2 CONH 2 +2H 2 O → (NH 4 ) 2 CO 3 → 2NH 3 + H 2 O+ CO 2 On a small scale ammonia is obtained from ammonium salts which decompose when treated with caustic soda or lime. 2NH4Cl + Ca(OH)2 → 2NH3 + 2H2O + CaCl2 (NH4)2 SO4 + 2NaOH → 2NH3 + 2H2O + Na2SO4 On a large scale, ammonia is manufactured by Haber’s process. N2(g) + 3H2(g) → 2NH3(g); ΔH = – 46.1 kJ mol-1

8 Flow chart for the manufacture of ammonia

9 Phosphine Phosphine is prepared by the reaction of calcium phosphide with water or dilute HCl. Ca3P2 + 6H2O → 3Ca(OH)2 + 2PH3 Ca3P2 + 6HCl → 3CaCl2 + 2PH3 In the laboratory, it is prepared by heating white phosphorus with concentrated NaOH solution in an inert atmosphere of CO2. Р 4 + 3КОН + 3Н 2 О → РН 3 + 3КН 2 РО 4 PH4I (phosphonium iodide)+KOH→KI + H2O+PH3

10 Properties of Ammonia Ammonia gas is highly soluble in water. Its aqueous solution is weakly basic due to the formation of OH– ions. NH 3 (g) + H 2 O(l) → NH 4 + (aq) + OH – (aq) 3CuO + 2NH 3 → 3Cu + 3H 2 O + N 2 It forms ammonium salts with acids, e.g., NH 4 Cl, (NH 4 ) 2 SO 4, etc. As a weak base, it precipitates the hydroxides of many metals from their salt solutions. 2FeCl 3 +3NH 4 OH→Fe 2 O 3 xH 2 O (brown ppt)+ 3NH 4 Cl The ammonia molecule can act is a Lewis base Cu 2+ (aq, blue) + 4 NH 3 (aq, deep blue) → [Cu(NH 3 ) 4 ] 2+ (aq) Ag + (aq, colourless) Cl - → AgCl (s, white ppt) AgCl (s, white ppt)+2NH 3 (aq) →[Ag(NH 3 ) 2 ]Cl)(aq,colourless)

11 Properties of Ammonia Ammonia Fountain. Demonstration of the high solubility of gaseous ammonia in water Preparation of Potassium Amide. Potassium amide is prepared by dissolution of potassium in liquid ammonia.

12 Nitrogen oxides


14 Nitric acid In the laboratory, nitric acid is prepared by heating KNO3 or NaNO3 and concentrated H2SO4 in a glass retort: NaNO3 + H2SO4 → NaHSO4+ HNO3

15 Preparation of Nitric acid

16 Nitric acid In aqueous solution, nitric acid behaves as a strong acid giving hydronium and nitrate ions. HNO 3 (aq) + H 2 O(l) → H 3 O + (aq) + NO 3 – (aq) Concentrated nitric acid is a strong oxidising agent and attacks most metals except noble metals such as gold and platinum. 8HNO 3 (dilute)+ 3Cu → 3Cu(NO 3 ) 2 + 2NO + 4H 2 O 4Sn + 10HNO 3  4Sn(NO 3 ) 2 + NH 4 NO 3 + 3H 2 O

17 Some metals (e.g., Cr, Al) do not dissolve in concentrated nitric acid because of the formation of a passive film of oxide on the surface. Concentrated HNO3 also oxidises non–metals and their compounds. Iodine is oxidised to iodic acid, carbon to carbon dioxide, sulphur to H2SO4, and phosphorus to phosphoric acid. I2 + 10HNO3 → 2HIO NO2 + 4H2O C + 4HNO3 → CO2 + 2H2O + 4NO2 S8 + 48HNO3(conc.) → 8H2SO4 + 48NO2 + 16H2O P4 + 20HNO3(conc.) → 4H3PO NO2 + 4H2O

18 Brown Ring Test The familiar brown ring test for nitrates depends on the ability of Fe2+ to reduce nitrates to nitric oxide, which reacts with Fe2+ to form a brown coloured complex.

19 Devarda's test Devarda's alloy (Cu/Al/Zn) is a reducing agent. When reacted with nitrate in sodium hydroxide solution, ammonia is liberated. 3 NO−3 + 8 Al + 5 OH − + 18 H 2 O → 3 NH [Al(OH) 4 ] − Aluminium is the reductant in this reaction.

20 Test for Ammonia Test for Ammonia using Nessler's Agent. Ammonia is tested in a 1:10 dilution row using K2[HgI4].

21 Phosphorus. Allotropic Forms White phosphorus is a translucent white waxy solid. It is poisonous, insoluble in water but soluble in carbon disulphide and has chemiluminescence. Р 4 + 3NaOH + 3H2O → PH3 + NaH2PO2 Р 4 + 3О 2 → 2Р 2 О 3 Р 2 О 3 + О 2 → Р 2 О 5 3Mg + 1/2P 4 → Mg 3 P 2 2P + 3Cl 2 → 2PCl 3

22 White phosphorus White phosphorus exposed to air glows in the darkness

23 Red phosphorus It is obtained by heating white phosphorus at 573K in an inert atmosphere for several days. When red phosphorus is heated under high pressure, a series of phases of black phosphorus are formed. Red phosphorus possesses iron grey lustre. It is odourless, non-poisonous and insoluble in water as well as in carbon disulphide. Chemically, red phosphorus is much less reactive than white phosphorus. It does not glow in the dark.

24 Black phosphorus It has two forms α-black phosphorus and β-black phosphorus. α-Black phosphorus is formed when red phosphorus is heated in a sealed tube at 803K. It can be sublimed in air and has opaque monoclinic or rhombohedral crystals. It does not oxidise in air. β-Black phosphorus is prepared by heating white phosphorus at 473 K under high pressure. It does not burn in air up to 673 K.

25 Phosphorus Tri- and Pentachlorides They are obtained by the action of thionyl chloride with white phosphorus. P4 + 8SOCl2 → 4PCl3 + 4SO2 + 2S2Cl2 P4 + 10SOCl2 → 4PCl5 + 10SO2

26 PCl 3 + 3H 2 O → H 3 PO 3 + 3HCl PCl 5 + H 2 O → POCl 3 + 2HCl POCl 3 + 3H 2 O → H 3 PO 4 + 3HCl CH 3 COOH + PCl 5 → POCl 3 + HCl + CH 3 COCl H 2 SO 4 + PCl 5 → POCl 3 + HCl + SO 2 (OH)Cl POCl 3 + 3H 2 O → H 3 PO 4 + 3HCl POCl 3 + 3ROH → PO(OH) 3 + 3RCl 2Ag + PCl5 → 2AgCl + PCl3 Sn + 2PCl5 → SnCl4 + 2PCl3 Reaction of Phosphorus Halides

27 Oxoacids of Phosphorus

28 Properties of phosphorus oxoacids and their salts The acids which contain P–H bond have strong reducing properties. 4 AgNO3 + 2H2O + H3PO2 → 4Ag + 4HNO3 + H3PO4 Acids in +3 oxidation state of phosphorus tend to disproportionate to higher and lower oxidation states. 4H3PO3 → 3H3PO4 + PH3 3H(PH 2 O 2 ) + 2HNO 3  3H 2 (PHO 3 ) + 2NO  + H 2 O 5H 4 P 2 O 6 + 3H 2 SO 4 + 2KMnO 4 + 2H 2 O  10H 3 PO 4 + 2MnSO 4 + K 2 SO 4 Salts of phosphorus oxoacids hydrolyze and base or neutral medium occurs

29 Applications of nitrogen compounds As a modified atmosphere, pure or mixed with carbon dioxide, to preserve the freshness of packaged or bulk foods Nitrogen can be used instead of CO2 to pressurize kegs of some beers, in particular, stouts and British ales, due to the smaller bubbles it produces, which make the dispensed beer smoother and headier Liquid nitrogen is used in the cryopreservation of blood, reproductive cells (sperm and egg), and other biological samples. It is used in the clinical setting in cryotherapy to remove cysts and warts on the skin.

30 Applications of nitrogen compounds Nitrous oxide (N 2 O), "laughing gas“, was discovered early in the 19th century to be a partial anesthetic, though it was not used as a surgical anesthetic until later. Nitrogen-containing drugs are drugs derived from plant alkaloids, such as morphine (there exist many alkaloids known to have pharmacological effects; in some cases, they appear natural chemical defenses of plants against predation). Drugs that contain nitrogen include all major classes of antibiotics and organic nitrate drugs like nitroglycerin and nitroprusside that regulate blood pressure and heart action by mimicking the action of nitric oxide.

31 Applications of phosphorous compounds White phosphorus, called "WP" (slang term "Willie Peter") is used in military applications as incendiary bombs, for smoke-screening as smoke pots and smoke bombs, and in tracer ammunition. The spontaneous combustion of phosphine is technically used in Holme’s signals. Containers containing calcium carbide and calcium phosphide are pierced and thrown in the sea when the gases evolved burn and serve as a signal. Phosphine is also used in smoke screens.

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