We studied the various isotopes of hydrogen or as well the two forms of the hydrogen molecule. In this unit we shall be learning the ways hydrogen is manufactured, its properties or uses.
Manufacture of Hydrogen:
Most if not all the process used to manufacture hydrogen make utilize of water. Water is a normal abundant source for the manufacture of hydrogen. Water can be reduced to hydrogen either chemically and electrically.
(a) Manufacture through chemical means:
The general reducing agents are Coke, (obtained from the destructive distillation of coal), carbon monoxide or hydrocarbon.
(i) The hydrogen is manufactured by allowing steam to react with red hot coke at about 1250k
C + H2O +1250K → CO + H2 + Water gas
ii) The mixture of CO and H2, recognized as water gas, is as well called synthesis gas since it is used for the synthesis of methanol or other hydrocarbons.
(iii) Hydrogen is as well produced by the reaction of gas (chiefly methane) with steam in the presence of nickel catalyst.
CH4 + H2O N1 catalyst → CO + 3H2
920 - 1170K
Similar reaction can occur with other hydrocarbons. In both cases above CO is converted to CO2.
[CO + H2O Catalyst → CO2 + H2]
450 - 650K
Hydrogen or carbon dioxide can easily be divided from each other via bubbling the gas mixture through water in which CO2 is quite soluble or H2 is virtually insoluble. In the presence of catalysts (like silica and alumina), at higher temperatures, the hydrocarbon break up and rearrange in what are termed as cracking reactions. Such reactions, which are utilized in refining of petroleum, produce hydrogen as a bi-product. One instance of simple cracking reaction is cracking of propane.
2 CH3 - CH2 - CH3 Catalyst + 875K → CH4 + CH2 = CH2 + CH3 - CH = CH2 + H2
Or 2C3H8 +catalyst+875K → CH4+C2H4+C3H6+H2
(b) Manufacturing by electrolysis:
(i) Electrolysis of acidified water using platinum electrodes is a convenient source of hydrogen (and oxygen). On a large scale, very pure hydrogen (>99.95%) can be attained from the electrolysis of aqueous solution of barium hydroxide between nickel electrodes. Hydrogen obtained via electrolysis of water is relatively expensive since of the cost of electrical energy.
(ii) Hydrogen can though be obtained cost-effectively as a bi- product in the electrolysis of brine during the manufacture of sodium hydroxide. During electrolysis, there is a competition at the anode between the oxidation of chloride ion or the oxidation of water.
For example: 2Cl- aq → Cl2(g) + 2e-
2H2O → O2 (g) + 4 H + aq + 4e-
When a concentrated salt solution (brine) is utilized, the first reaction preferentially takes place at the cathode; the reaction is the reduction of water since it is more easily decreased than Na+
2H2O (1) +2e H2 (g) + 2OH- (ag)
The anode and the cathode reactions are combined to give the reactions thus:-
2 Na+ + 2C1- + 2H2O → Cl2 (g) + H2 (g) + 2Na+ + 20H- + aq
In the laboratory hydrogen can be prepared by the reaction of water or dilute acids on electropositive metals these as alkali metals alkaline earth metals, the metals of group 12 (example Zn) and the lanthanides. The reaction can be explosively violent with alkali metals (example II , Rb) convenient laboratory process employ sodium amalgam and calcium with water or zinc and tin with hydrochloric acid.
Properties of Hydrogen:
(i) Hydrogen is the lightest element known. It is colorless, odorless and tasteless gas. The hydrogen molecule is
(ii) Thermally stable and has little tendency to dissociate at normal temperatures, the reaction H2g → 2H (g) = +436kg mold-1 being an endothermic one. Though, at high temperature in an electric and under ultraviolet light, H2 does dissociate. The atomic hydrogen produced, exists for less than half a second after that it recombines to give molecular hydrogen and liberates a large amount of energy (436kg mol-1 ), in form of heat. Most of the change metals catalyse the combination reaction of hydrogen.
(iii) Atomic hydrogen is a powerful reducing agent and reduces copper, silver and mercury salts to the metallic state of.
For example: Ag NO3 + H → Ag + HNO3
(iv) It merges by alkali metals to form hydrides
For example: Na+ H → NaH
(v) It reduces sulphur to hydrogen sulphide
For example S+2H → H2S
(vi) Carbon monoxide is reduced to formaldehyde.
For example CO +2H →H CHO
(vii) It also reacts with oxygen at room temperature to form hydrogen peroxide.
For example O2 ± 2H →H2O2
(viii) H is employed for wielding. Atomic hydrogen is produced via passing ordinary hydrogen by electric arc maintained between 2 electrodes. The atoms set free are carried away via a stream of incoming hydrogen gas. Such free atoms recombine at once on coming in contact through a metallic surface liberating a large amount of heat and therefore arising temperature of the metal to say 4000 - 5000k. This principle is used in the making of the 'atomic hydrogen welding torch'. It provides a chance of welding at an extremely high temperature but in a reducing environment.
(ix) Although the fairly high bond dissociation energy of the hydrogen molecules it is moderately reactive and forms strong bonds by many other elements. It reacts through almost all elements except the noble gases.
(x) Hydrogen reacts by alkali and alkali earth metals via accepting an electron to form conic hydrides; for example KH, CaH2
(xi) By non-metals it forms covalent hydrides, for example NH3, H2O and HF.
Fig: Atomic hydrogen welding torch
(xiii) Hydrogen is effortlessly oxidized to water and; therefore it acts as a very good reducing agent in a variety of situations
(xii) Hydrogen is used in metallurgy to reduce metal oxides to metals in cases where carbon cannot be used because the metal can form carbide. Such metals include Mo and W.
WO + 3H2 → W + 3H20
(NH4) MoO4 + 3H2 → n Mo 2NH3 + 4H20
(xv) Hydrogen adds on the multiple bonds in organic compounds. In the presence of catalysts these as finely divided nickel, palladium or mixtures of metal oxides, unsaturated organic compound are therefore reduced to saturated compounds. For example.
CH3CH = CH CH3 + H2 → CH3 CH2 CH2 CH3
2 - Butane Pd, 1 atm n- butane
CH3C = N + 2H2 → CH3CH2NH2
Methyl cyanide ZnO/Cr2O3 ethylamine
Catalytic hydrogenation of unsaturated liquid vegetable oils to solid edible fats demonstrates the industrial application of the reduction reactions; for instance the reduction of an oleate (ester of oleic acid) to ester of stearic acid.
CH3 (CH2)7 CH = CH (CH2)7 COOR + H2 → CH3 (CH2)16 COOR
Hydrogen reacts by carbon monoxide in the occurrence of catalysts to form methanol.
CO + 2H2 → CH3OH
This reaction is recognized as hydroformylation reaction and is utilized in the industrial preparation of methanol.
Uses of Hydrogen:
Several of the utilizes of hydrogen contain the following:
1. The largest single employ of hydrogen is in the syntheses of ammonia that is employed in the manufacture of nitric acid and nitrogenous fertilizers.
2. Hydrogen is utilized in the hydrogenation of vegetable of oils and the manufacture of methanol
3. In space crafts, hydrogen gas is used in fuel cells for generating electrical energy and for providing clean drinking water to the astronauts. In a fuel cell, electrical energy is produced via the reaction of hydrogen leaf. This is sometimes termed "cold combustion". A hydrogen oxygen fuel cell might be having an alkaline of acidic electrolyte.
The one in fig has porous carbon electrodes and KOH as electrolyte.
Fig: A hydrogen-oxygen fuel cell with KOH electrolyte and porous carbon electrodes
The half cell reactions are given below:
H2 (g) + 20H- → 2H2O + 2e E0 - 0.8280 anode
1 /2 O2 (g) + H2O (1) + e → 2OH- E° = 0.401V, cathode
H2 (g) +1/2 O2(g) → H2 0(1) E° = 1.229V
By acidic electrolyte, the half-cell reacts irons are:
H2 (g) → 2H + 2e, E0 = 0 Anode
1/2 O2 (g ) + 2H+ +2e → H2O(l) E0 = 1.22911 cathode
H2 (g) + 1/2 O2 (g) → H2O (l), E0 = 1.229v
We can see from the equations above that the electromotive force of the cell remains similar whether we employ alkaline or acidic electrolyte. This is because we are using the same reactants at the electrodes in both cases.
Fuel cells have numerous advantages over other sources of energy. Initially in a fuel cell unlike in the dry cell or storage battery (that requires recharging as well), the cathode and anode reactants are continuously supplied so that energy can be indefinitely withdrawn from it. Secondly, in a fuel cell energy is extracted from the reactants under almost ideal conditions. Therefore, the thermo dynamic efficiency of the fuel cells is higher than of mainly of the ordinary combustion processed. Fuel cells have efficiencies approaching 75% whereas power plants that burn fuels have efficiencies of only about 40%.
Combustion of hydrogen is a highly exothermic reaction and produces no pollutants;
2 H2 (g) + O2 (g) → 2H2O (j), ΔH = -572kg
Liquid hydrogen is therefore, employed as a rocket fuel. It is now clear that world reserves of fossil fuels like coal, oil and gas are finite, so they cannot last forever. Nuclear and hydro electric power cannot meet the world's energy needs. Furthermore, these resources post a danger to the world's environment. With such facts in mind there is now an active search for alternative source of energy. In addition to solar power, hydrogen is being believed a potential fuel for the future.
Hydrogen as a fuel has many benefits over the conventional fossil fuels and electric power. It is available in unlimited quantities in sea water. It is pollution free because the major product of its combustion is water by only traces of nitrogen oxides. It liberates greater energy per limit weight of fuel in comparison to gasoline and other fuels.
Hydrogen can be transported as a gas in high pressure pipelines, as a liquid in tankers and even as a solid in form of metal hydrides. Unlike electricity hydrogen can be stored and utilized when required. Hydrogen though has the following disadvantages via: Hydrogen like electricity is a secondary source of energy since it is produced using energy from a primary source these as coal, nuclear fission or sun. Preparation of hydrogen through electrolysis is not economical at present in actuality more energy has to be spent in electrolysis of water than what can be released via burning hydrogen as a fuel. Decomposition of water via solar energy in presences of catalysis is recognized as photochemical decomposition of water. Using catalysts scientists in France have been able to achieve the proficient decomposition of water under visible and ultra violet light. If this procedure can be made industrial, a convenient technique of converting solar energy directly to a useful form of store chemical energy will be accessible.
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