Hydrogen

1st element of periodic Table, position uncertain, Atomic No. 1, Atomic wt.1.008, Nucleus with 1 Proton, No Neutron, 1 planetary Electron present in s- orbital of the k - Shell or 1st orbit. Hydrogen atom has the tendency to loose electron and become H+ ion (Electropositive character) like Alkali metals. The atom also has tendency to gain electron & become H -ion (Electronegative Character) like halogens. H2 does not occur free in earth's atmosphere however it forms 90% of sun's atmosphere. Except for traces i.e., 1 part in 1.5 million parts Water containg small amount of an acid or base when electrolysed gives H2 at cathode and O2 at anode. Case I - Suppose the hydrolysis is carried out in acidic solution (H2SO4 for example ) than at Anode , Sulphate ions (SO4- -) (although present in much larger quantity ) are not discharged as their DISCHARGED POTENTIAL is very much higher than that of the hydroxyl ions . Case II - Similarly suppose hydrolysis is carried out in an alkali solution (KOH for example) than at cathode Hydrogen ions are discharged & not Potassium ions as the discharge potential of K+ is very much higher. By the action of certain active metals on water - Alkali metals are most active for this reaction. Large amount of heat is given off in most of the reaction and very often H2 evolved catches fire. Note: Reaction between calcium and water Rapid and nonviolent. Reaction with magnesium when water is warmed only. Iron and zinc are even less reactive. They decompose water only when superheated steam is led over them at high temp. Most other elements are far less reactive to be able to displace H2 from water . By the action of water on a metal hydride - Especially hydrides of alkali and alkaline earth metals Metals like Zn, Sn & Al (present below Hydrogen in the electromotive series) have strong tendency to pass into acidic solutions as positive ions . In doing so they lose electrons, which are gained by positively charged Hydrogen ions and change into neutral Hydrogen atoms. The atoms unite to form molecules. Certain metals like Al, Zn and Sn react with alkali Evolving H2 2Al + 2NaOH + 2H2O 2NaAlO2 + 3H2 Zn + 2NaOH Na2 ZnO2+ H2 Sn + 2NaOH + 2H2O Na2SnO2 + H2 Pure Hydrogen is produced by electrolysis of warm conc. Solution of Ba(OH)2 in a hard glass tube using Ni electrodes. Hydrogen is liberated at cathode while oxygen is liberated at anode. Pure hydrogen can also be prepared by Bosch process in which reduction of water vapour (steam) takes place with Carbon. Colorless, Tasteless, Odorless Gas ; Lightest element known ; Density .0695 gm/ cc. ; CRITICAL TEMP. -236.9°, At 1 Atm. Pressure it can be liquefied at -252.8° and solidified at -259.8°c. DIATOMIC (CP / Cv = 1.40) ; Metals like Fe ,Au, Pt, Pd , Ni etc can adsorb or occlude large volumes of H2 gas at different temperatures. (Pd can occlude 1000 times of its own volume ) Covalent compounds are formed on reaction with nonmetals like O2 , N2 , F2 , Cl2 , Br2 etc . (But not reactive at ordinary temperatures) At ordinary temperature Hydrogen exists in molecular state in which atoms have stable He structure. To be converted into atomic hydrogen either high temperature or a catalyst is required. H2 molecules are first adsorbed at the surface of catalyst and dissociated into atoms. These atoms than combine with Non Metals. With metals like Li, Na, K & Ca etc. hydrogen forms HYDRIDES. Hydrogen reduces oxides of less electropositive elements e.g. CuO + H2 Cu +H2O ; Fe3O4 + 4H2 3Fe + 4H2O . Note: However Oxides of alkali and alkaline earths are not reduced by H2. H2 atom has a tendency to change into hydrogen ion or proton (H+) by losing its solitary e- . The proton has a remarkable property of acting as an electron acceptor that's why it rarely exists in free state. E.g. It readily combines with H2O molecules to form [H3O]+ (HYDROXONIUM ION or Hydrated Hydrogen Ion.) Hence we can say that due to this tendency the following conversion can be easily done H+ + e- H ; H + H H2. Thus strongly electropositive metals (Na, K, Zn, Mg, Sn etc ) lying below H2 in electromotive series , liberate H2 when treated with a solution containing Protons (H+). INFACT THESE METALS LOOSE THEIR VALANCE e- To H+. However weak electropositive metals (Cu, Ag, Au, hg etc ) i.e., those lying above H2 in electromotive series does not liberate H2 on treatment with same solution (containing H+) as these DO NOT LIBERATE THEIR VALANCE e- for H+. Two different forms of H2 molecule i.e., ortho and para hydrogen are discovered. The nucleus of an atom has nuclear spin (In the same way as e- spin). In H2 molecule the two nuclei may be spinning either in same direction or in opposite directions. This give rise to SPIN ISOMERISM (also seen in D2, N2 , F2 , Cl2 etc). The form with spin in same direction is called ORTHO form. The form with spin in opposite direction is called PARA form. There are considerable differences seen between the physical properties (b.p., Specific Heats , Thermal Conductivity) of Ortho and Para forms because of difference in their internal energies . Para form has lower energy & at Absolute Zero the gas Contains 100 % of Para form. As temperature rises para form is converted in to ortho form. Ordinary Hydrogen is an equilibrium mixture of 75% ortho and 25% para Hydrogen. As both the forms are interconvertible at room temperature both of the forms are in equilibrium. On decreasing the temperature % of para Hydrogen increases in the equilibrium mixture. Ortho Hydrogen is more stable. Para Hydrogen has the tendency to change itself in to ortho Hydrogen. Para Hydrogen is usually prepared by passing a mixture of the two forms through a tube packed with charcoal cooled to liquid temperature. Para Hydrogen prepared in this way is kept for weeks at room temperature in a glass vessel, because the ortho-para conversion is slow in the absence of catalysts. Suitable catalyst include activated charcoal, atomic Hydrogen, and metals such as Fe, Ni, Pt, W etc., along with paramagnetic substances or ions e.g., O2, NO, NO2, Co2+, Cr2O3. Thenard, a French scientist, prepared hydrogen peroxide in 1818 by the action of dilute hydrochloric acid on barium peroxide (BaO2). He called it as OXYGENATED WATER. Other methods for the preparation of H2O2 are Na2O2 is added slowly in ice cold 20% H2SO4. On cooling crystals of Na2SO4.10H2O separate out. These are filtered off. Some impurity of Na2SO4 remains present in H2O2. Powered barium peroxide is mixed in water and then CO2 is passed in it. By the action of dilute H2SO4 on hydrated barium peroxide: In this method anhydrous barium peroxide (BaO2) is not used because a layer of BaSO4 is deposited on it. This prevents further reaction. But this difficulty does not come when hydrated barium peroxide is used. H2SO4 catalyses the decomposition of H2O2 that's why some weak acid such as H3PO4, H2CO3 etc., is preferred. Electrolysis of 50% H2SO44 is done at 0°C using platinum electrodes, Perdisulphuric acid (H2S2O8) is formed. Perdisulphuric acid on hydrolysis with water gives H2O2. H2O2 formed is separated by distillation under reduced pressure. When (NH4)2SO4 and H2SO4 are mixed in 1: 1 molar ratio, Ammonium hydrogen sulphate is formed. Which is electrolyzed to give Ammonium persulphate by using Pt electrodes at 00 C. Ammonium persulphate formed is hydrolyzed by heating with water at low pressure (43 mm Hg). H2O2 is produced. H2O2 is now manufactured by a cyclic process in which 2-Ethyl or 2-Butyl anthroquinol is oxidized by air to the corresponding quinone and H2O2. The anthroquinone thus produced is reduced back to anthroquinol at a moderate temperature by using Pt, Pd, or raney Ni catalyst. The reaction is carried out in a mixture of organic solvents [ester/ hydrocarbon or octanol/ methyl naphthalene]. The solvent must : Dissolves quinol and quinone. Resist oxidation. Be immiscible with water. Pure anhydrous H2O2 is light blue, odorless syrupy liquid. It produces blisters on skin. It is soluble in water, alcohol and ether. Its b.p. is 152°C and freezing point is - 0.89°C. Hydrogen peroxide is associated through inter molecular hydrogen bonds, therefore its b.p. is higher than expected. Its density at 0°C is 1.45 g/cc. Dielectric constant of H2O2 is 93.7. Addition of water increases the dielectric constant of solution. H2O2 molecule is diamagnetic. H2O2 molecule is polar and its dipole moment is 2.1 D. It has both polar and non-polar bonds. H2O2 decomposes on heating. It is an unstable liquid. Finely divided Fe, Cu, Au, Pt, MnO2, metals ions Cu++, Ni++ and traces of base act as positive catalyst and their presence increases the decomposition. H3PO4, C6H5NHCOCH3 (acetanilide), C2H5OH, glycerol etc. act as negative catalyst and slow down the decomposition of H2O2. H2O2 is acidic in nature but its aqueous solution is neutral towards litmus. It dissociates in water to give hydronium ion . H2O2 + H2O = H3O+ + HO-2 (Ka = 1.4 x 10-12 at 298 K) Value of dissociation constant (Ka) suggests that its aqueous solution is more acidic than water. For water Kw = 1 x 10-14 at 25 °C (298K) Alkalis and carbonates react with H2O2 forming corresponding peroxides. In these reaction hydrogen atoms of H2O2 are replaced by metal ions which confirm the acidic nature of H2O2. H2O2 is stronger acid than H2CO3 H2O2 is a strong oxidizing agent. It accepts the electron and itself is reduced in acid or alkaline medium. H2O2 + 2H+ + 2e- 2H2O (E°= + 1.77V) HO-2 + H2O + 2e- 3OH - (E°= +0.87V) Thus in acid medium H2O2 is reduced into H2O and in alkaline medium into OH-. In both these reactions oxidation number of oxygen increases from -1(H2O2) to -2(H2O). In alkaline solution reactions are generally fast. Examples of oxidizing reaction are: (a) Lead sulphide (PbS) is oxidized into PbSO4. KNO2 is oxidized into KNO3 Na2SO3 is oxidized into Na2SO4. (d) Sodium arsenite (Na3AsO3) is oxidised into sodium arsenate. (e) Hydrogen sulphide is oxidised into sulphur. (f) Ferrous sulphate is oxidised into ferric sulphate. (g) KI is oxidised into I2. (h) Formaldehyde (HCHO) is oxidised into formic acid (i) Potassium ferrocyanide is oxidised into potassium ferricyanide. (j) In presence of FeSO4 benzene is oxidised into phenol. (k) Hg is oxidised into HgO Cold H2O2 is added in cold mixture of K2Cr2O7 and concentrated H2SO4. A blue solution of blue perchromate (CrO5) is obtained. Note: Oxidation No. Of Cr in CrO5 is +6 and not +10 H2O2 is electron donor and acts as reductant. Reactions are fast in alkaline medium Examples of reduction reactions of H2O2 are Silver oxide (Ag2O) is reduced into Ag Lead dioxide (PbO2) is reduced into PbO (c) Cl2 is reduced into HCl (d) Gold oxide is reduced into gold (e) In acid medium MnO2 is reduced into MnSO4 (f) Acidified KMnO4 is reduced by H2O2. Pink color of KMnO4 disappears. (g) In alkaline medium potassium ferricyanide is reduced into potassium ferrocyanide. H2O2 and O3 both reduces each other Bleaching action of H2O2 is due to its oxidizing properties. It bleaches silk, cotton, wool, hair, feather, ivory etc. Uses of H2O2 H2O2 is germicide and antiseptic. It is used for washing cuts, wounds, teeth, ears etc. and in gargles under the name perhydrol. It prevents the putrification of milk, liquor etc. Its dilute solution is used in the bleaching of cotton, wool, silk, hair, ivory, paper, pulp etc. Old oil paintings turn black with time. H2O2 is used to restore their color. Paintings turn black due to the formation of black PbS. Hydrogen peroxide oxidizes black PbS into white PbSO4. It is used as an oxidizing agent. It is used in the preparation organic and inorganic compounds such as Na2O2, sodium perborate, epoxides, peracids such as H2SO5 (permonosulphuric acid, caro's acid), H2S2O8 (perdisulphuric acid, Marshall's acid) etc. It is used as a laboratory reagent in the test of Ti4+ , Cr3+ and V5+ ions. It is used as propellant for rockets, torpedoes etc and as a fuel. As a propellant it is used to oxidize alcohol, petrol, hydrazine etc. N2H4 + 2H2O2 N2 + 4H2O H2O2 is used as antichlor to remove Cl2 etc present in a solution. Strength of solution of H2O2 Strength of H2O2 solution may be represented in several ways such as grams/lit, moles/litre, normality, ppm (mg/litre), volume etc. Concentration in Volume It is equal to the volume of O2 obtained at STP by decomposing one volume of H2O2 solution. If the strength of H2O2 solution is 20 volume, this means that 1cc H2O2 solution on decomposition gives 20CC O2 at STP or one litre H2O2 solution on heating gives 20 litre O2 at STP. Concentration in percentage3 It is equal to the grams of H2O2 present in 100 ml of H2O2 solution. Thus if the strength of a H2O2 solution is 8%, this means that 8 grams of H2O2 are present in its 100 ml solution.