JEE Chemistry Hydrogen Notes Set 03

Position and Isotopes of Hydrogen

1. Hydrogen (1s¹) resembles alkali metals: one valence electron, forms H⁺, valency 1, a good reducing agent and forms a stable oxide H₂O. It also resembles halogens: diatomic molecule, forms H^- with strongly electropositive metals, one electron short of the He configuration.
2. Because of this dual behaviour it has no fixed position and is usually placed at the top of group 1. Its ionisation energy (1312 kJ/mol) is close to that of halogens and much higher than that of alkali metals.
3. Isotopes: protium (¹H, no neutron, 99.985%), deuterium (²H or D, one neutron) and tritium (³H or T, two neutrons, radioactive, beta emitter with half-life about 12.3 years). They have the same electronic configuration and chemical behaviour, but differ in rate of reaction and physical properties.
4. Three isotopes give six diatomic molecules: H₂, D₂, T₂, HD, HT and DT. Hydrogen is the most abundant element in the universe, and the ninth most abundant element in the earth's crust.
5. Dihydrogen occurs as nuclear-spin isomers: ortho-hydrogen (parallel spins, higher energy) and para-hydrogen (antiparallel spins, lower energy). Low temperature favours para-hydrogen, and at room temperature the ratio is about 75% ortho and 25% para.

Preparation of Dihydrogen

1. Laboratory method: zinc (or iron) with dilute HCl or dilute H₂SO₄: Zn + 2HCl gives ZnCl₂ + H₂. Metals above hydrogen in the electrochemical series displace it from dilute non-oxidising acids. Dilute HNO₃ does not give H₂ with most metals (except Mg and Mn in cold dilute HNO₃).
2. Amphoteric metals and non-metals react with alkalis: Zn + 2NaOH gives Na₂ZnO₂ + H₂; 2Al + 2NaOH + 2H₂O gives 2NaAlO₂ + 3H₂; Si + 2NaOH + H₂O gives Na₂SiO₃ + 2H₂.
3. Active metals react with water: 2Na + 2H₂O gives 2NaOH + H₂ (cold). Mg, Al and Zn react with hot water or steam, and Fe reacts with red-hot steam: 3Fe + 4H₂O gives Fe₃O₄ + 4H₂.
4. Hydrides react with water: NaH + H₂O gives NaOH + H₂ and CaH₂ + 2H₂O gives Ca(OH)₂ + 2H₂.
5. Industrial methods: electrolysis of acidified or alkaline (Ba(OH)₂) water gives very pure hydrogen; steam reforming of hydrocarbons CH₄ + H₂O gives CO + 3H₂ over a nickel catalyst at about 1270 K; water-gas shift CO + H₂O gives CO₂ + H₂ (Bosch process, using water gas). Hydrogen is also a by-product of the electrolysis of brine.

Properties and Uses of Dihydrogen

1. Dihydrogen is colourless, odourless, tasteless, the lightest gas, and only slightly soluble in water. It burns with a pale blue flame and forms explosive mixtures with air or oxygen. The H-H bond energy is high (436 kJ/mol), so it is inert at room temperature.
2. It acts as a reducing agent: CuO + H₂ gives Cu + H₂O, Fe₃O₄ reduced to Fe. It cannot reduce oxides of active metals such as Al₂O₃, MgO or CaO. It reacts with non-metals: N₂ (Haber process), Cl₂, O₂, S.
3. Hydrogenation: unsaturated vegetable oils are converted to saturated fats with H₂ over a nickel catalyst (vanaspati ghee).
4. Atomic hydrogen (produced in an electric arc) is much more reactive; it is used in atomic hydrogen welding. Atomic hydrogen reacts with oxygen to give mainly H₂O₂.
5. Uses: synthesis of NH₃ and methanol, hydrogenation of oils, rocket fuel (liquid hydrogen), fuel cells, welding and metallurgy as a reducing atmosphere.

Hydrides

1. Ionic (saline) hydrides: formed by s-block metals (except Be and Mg which are largely covalent). They contain H^-, are white crystalline solids with high melting points, and react violently with water giving H₂. On electrolysis of the molten hydride, hydrogen is liberated at the anode. Fires caused by them are put out with dry sand, not water or CO₂.
2. Covalent (molecular) hydrides: formed by p-block elements. They may be electron-deficient (B₂H₆), electron-precise (CH₄, SiH₄) or electron-rich (NH₃, H₂O, HF). Boiling point rises down a group except for NH₃, H₂O and HF, which are high due to hydrogen bonding.
3. Metallic (interstitial) hydrides: formed by d- and f-block metals (for example TiH_1.73, PdH_0.8). They are non-stoichiometric, conduct electricity and are less dense than the parent metals. Group 7, 8 and 9 metals (except Pd) do not form hydrides, called the hydride gap.
4. Oxidation state of hydrogen is +1 in most compounds, -1 in metal hydrides such as CaH₂, and 0 in H₂.

Water and Hardness

1. Water is a bent molecule with a high boiling point because of hydrogen bonding. Ice is less dense than water because of an open hydrogen-bonded cage, and water has maximum density at 4 degrees C. Water is a universal solvent because of its high dielectric constant.
2. Hard water does not lather with soap because Ca²⁺ and Mg²⁺ ions form insoluble calcium and magnesium salts of fatty acids. Temporary hardness is due to bicarbonates of Ca and Mg. Permanent hardness is due to chlorides and sulphates of Ca and Mg.
3. Removal of temporary hardness: boiling (Ca(HCO₃)₂ gives CaCO₃ + H₂O + CO₂; Mg(HCO₃)₂ gives Mg(OH)₂ precipitate); Clark's method (adding slaked lime, Ca(OH)₂).
4. Removal of permanent hardness: washing soda (Na₂CO₃), Calgon process (sodium hexametaphosphate, Na₆P₆O₁₈, forms soluble complexes with Ca²⁺ and Mg²⁺), permutit or zeolite process (Na⁺ exchanged for Ca²⁺ and Mg²⁺) and synthetic ion-exchange resins (cation exchange removes Ca²⁺ and Mg²⁺, anion exchange removes anions, giving demineralised water).
5. Synthetic detergents work in hard water because their calcium and magnesium salts are soluble.
6. Heavy water D₂O: obtained by repeated electrolysis of water. Freezing point 3.8 degrees C, maximum density at 11.6 degrees C. It is used as a moderator in nuclear reactors (slows fast neutrons) and as a tracer. Ionic compounds are less soluble in D₂O, and reactions of D₂O are slower than those of H₂O.

Hydrogen Peroxide

1. Preparation: BaO₂.8H₂O + H₂SO₄ (cold, dilute) gives H₂O₂ + BaSO₄; Na₂O₂ + H₂SO₄ (dilute, cold) gives H₂O₂ + Na₂SO₄; industrially by auto-oxidation of 2-ethylanthraquinol, or electrolysis of 50% H₂SO₄ followed by hydrolysis of peroxodisulphuric acid.
2. Structure: non-planar, open-book shape, with O-O single bond, two O-H bonds in different planes. The dihedral angle is about 111.5 degrees in the gas phase.
3. Acidic nature: weakly acidic. As an oxidising agent: H₂S + H₂O₂ gives S + 2H₂O; PbS + 4H₂O₂ gives PbSO₄ + 4H₂O; 2KI + H₂O₂ gives I₂ + 2KOH; Fe²⁺ goes to Fe³⁺ (Fenton's reagent is FeSO₄ with H₂O₂).
4. As a reducing agent (with stronger oxidants): 2KMnO₄ + 3H₂SO₄ + 5H₂O₂ gives K₂SO₄ + 2MnSO₄ + 8H₂O + 5O₂; Cl₂ + H₂O₂ gives 2HCl + O₂; it reduces Ag_{2O to Ag, O}3 _{to O}2 _{and acidic K}2_Cr2_O7~ to Cr³⁺.
5. Decomposition: 2H₂O₂ gives 2H₂O + O₂, speeded up by light, metal powders (Pt, Au), MnO₂, and slowed down by glycerol and acetanilide, which act as stabilisers (negative catalysts).
6. Strength: x volume H₂O₂ means 1 mL gives x mL of O₂ at STP. Volume strength = 5.6 x normality = 11.2 x molarity. Percentage (w/v) strength = volume strength x 0.3036. Example: 20 volume solution is about 6.07% and has normality 3.57.
7. Uses: bleaching of delicate fibres, antiseptic, rocket propellant, restoring old paintings (PbS black to white PbSO₄) and as a mild oxidiser in pollution control.