HYDROGEN

  • The most abundant element in the universe.
  • The third most abundant on the surface of the globe.
  • In atomic form it consists of only one proton and one electron. However, in elemental form it exists as a diatomic (H2) molecule and is called dihydrogen.

POSITION OF HYDROGEN IN THE PERIODIC TABLE

Hydrogen is the first element in the periodic table because the elements in the periodic table are arranged according to their electronic configurations, However, its placement in the periodic table has been a subject of discussion because of its resemblance with alkali metals as well as with halogens.

Resemblance with alkali metals.

  • Electronic configuration. Both hydrogen and alkali metals have similar valency shell configuration (Hydrogen : 1s1 , Alkali metals : ns1)
  • Electropositive nature (Oxidation States). Both Hydrogen and alkali metals lose one electron to form unipositive ions g. H+, Na+

(Due to very high ionization enthalpy hydrogen does not possess metallic characteristics under normal  conditions.)

  • Combination with electronegative elements. Like alkali metals, hydrogen forms oxides, halides and sulphides with non metals. g. (H2O, Na2O), (HCl, KCl), (H2S, Na2S) etc.
  • Reducing nature. Both alkali metals and hydrogen are very good reducing agents.

Resemblance with Halogens

  • Electronic configuration. Both halogens (ns2np5) and hydrogen (1s1) are short by one electron to the corresponding noble gas configuration. g. H(He), F(Ne)
  • Oxidation States. Both Hydrogen and halogens gain one electron to form uninegative ion. g. H, F

Ionisation enthalpy. In terms of ionization enthalpy, hydrogen resembles more with halogens, Di H of Li is 520 kJ mol–1, F is 1680 kJ mol–1 and that of H is 1312 kJ mol–1.

  • Halogens as well hydrogen exist as diatomic molecules.
  • Salt Formation. Like halogens it can combine with metals to form salts. g. NaH, NaCl
  • Combination with non-metals. Like halogens it reacts with non-metals such as C, Si, Ge etc. to form covalent compounds. g. CH4, CCl4

(In terms of reactivity, Hydrogen  is very less reactive as compared to halogens)

Neither resembles with alkali metals nor halogens

  • Ionic Size : Size of (H+) is ~1.5 × 10–3 However normal atomic and ionic sizes are ~ 50 to 200 pm.
  • Existence as free ion : H+ does not exist freely and is always associated with other atoms or molecules.

e.g.  

Thus, Hydrogen is unique in behavior and is, therefore, best placed separately in the periodic table

DIHYDROGEN, H2

Occurrence

Universe : 70% of the total mass

  • Planets : Jupiter and Saturn consist mostly of hydrogen
  • Earth’s atmosphere : 0.15% by mass
  • Earth’s crust and the oceans : 15.4%

In the combined form besides in water, it occurs in plant and animal tissues, carbohydrates, proteins, hydrides including hydrocarbons and many other compounds.

Isotopes of Hydrogen

Hydrogen has three isotopes: Protium (n=0, p=1), Deuterium or D (n=1, p=1),  and Tritium  or T (n=2, p=1).

In the year 1934, an American scientist, Harold C. Urey, got Nobel Prize for separating hydrogen isotope of mass number 2 by physical methods.

Relative abundance (%) : Protium (Hydrogen) = 99.985 % , Deuterium = 0.0156 % &

Tritium = 10–15 ­­%.

Tritium is radioactive and emits low energy b particles (t½, 12.33 years).

Atomic and Physical Properties of Hydrogen

  • Relative atomic mass (g mol–1) : H­­­2 (2.016) < D2 (4.028) <T2 (6.032)
  • Melting point / K : H­­­2 < D2 < T2
  • Boiling point/ K : H­­­2 < D2 < T2
  • Density / gL–1 : H­­­2 < D2 < T2
  • Enthalpy of fusion/kJ mol–1: H­­­2 < D2
  • Enthalpy of vaporization/kJ mol–1 : H­­­2 < D2

Enthalpy of bond dissociation/kJ mol–1 at 298.2K : H­­­2 < D2

Internuclear distance/pm : H­­­2 = D2

Ionization enthalpy : Hydrogen (1312 kJ/mol),

Electron gain enthalpy/kJ mol–1 : Hydrogen (-73 kJ/mol)

Covalent radius/pm : Hydrogen (37 pm),

Ionic radius(H)/pm : Hydrogen (208 pm)

Due to same electronic configuration isotopes have almost the same chemical properties.

Isotopes have different rates of reactions due to their different enthalpy of bond dissociation.

Physical properties of isotopes are different due to their large mass differences.

PREPARATION OF DIHYDROGEN, H2

There are a number of methods for preparing dihydrogen from metals and metal hydrides.

Laboratory Preparation

  • It is usually prepared by the reaction of granulated zinc with dilute hydrochloric acid.

Zn + 2H+→Zn2+ + H2

  • It can also be prepared by the reaction of zinc with aqueous alkali.

Commercial Production

The commonly used processes are outlined below:

  • Electrolysis of acidified water using platinum electrodes gives hydrogen.
  • High purity (>99.95%) : Electrolysis of warm aqueous Ba(OH)2 solution between nickel electrodes produce high purity dihydrogen at cathode.
  • Electrolysis of brine solution [NaCl (aq.)] produces dihydrogen at cathode as per following reactions,

at anode: 2Cl(aq)→Cl2(g) + 2e at cathode: 2H2O(l) + 2e→H2(g) + 2OH(aq)

The overall reaction is

2Na+ (aq) + 2Cl(aq) + 2H2O(l)

Cl2(g) + H2(g) + 2Na+ (aq) + 2OH(aq)

  • Reaction of steam on hydrocarbons or coke at high temperatures in the presence of catalyst yields hydrogen.

e.g.,

Water gas or Syngas: Mixture of CO and H2 .

Coal Gasification : The process of producing ‘syngas’ from coal is called ‘coal gasification’.

Water-gas shift reaction : The production of dihydrogen can be increased by reacting carbon monoxide of syngas mixtures with steam in the presence of iron chromate as catalyst.

This is called water-gas shift reaction. Carbon dioxide is removed by scrubbing with sodium arsenite solution.

Sources of dihydrogen : Petro-chemicals ~77% , Coal~18%, Electrolysis~4% & Other sources ~1%.

PROPERTIES OF DIHYDROGEN

Physical Properties

  • Dihydrogen is a colour less, odour less, taste less, combustible gas.
  • It is lighter than air and insoluble in water.

Chemical Properties

  • The H–H bond dissociation enthalpy is the highest for a single bond between two atoms of any element.
  • The dissociation of dihydrogen into its atoms is only ~0.081% around 2000K which increases to 95.5% at 5000K.
  • It is relatively inert at room temperature due to the high H–H bond enthalpy.
  • The atomic hydrogen is produced at a high temperature in an electric arc or under ultraviolet radiations.
  • Itselectronic configuration is 1s1 hence it can (i) loose one electron to give H+, (ii) gain an electron to form H, and (iii) share one electron to form a single covalent bond.

Chemical Reactions of dihydrogen

Reaction with halogens: H2(g) + X2 (g)→2HX (g) (X = F, Cl, Br, I)

Order of reactivity : I2 < Br2 < Cl2 < F2

Fluorine reacts even in the dark, whereas iodine requires a catalyst.

Reaction with dioxygen :

Reaction with dinitrogen :

This is the method for the manufacture of ammonia by the Haber process.

Reaction with metals: With many metals it combines at a high temperature to yield the corresponding hydrides

H2(g) + 2M(g)→ 2MH(s);

where M is an alkali metal

Reactions with metal ions and metal oxides: It reduces some metal ions in aqueous solution and oxides of metals (less active than iron) into corresponding metals.

H2(g) + Pd2+ (aq)→Pd(s) + 2H+(aq)

yH2(g) + MxOy(s)→xM(s) + yH2O(l)

Reactions with organic compounds: It reacts with many organic compounds in the presence of catalysts to give useful hydrogenated products of commercial importance. For example :

(i)   Hydrogenation of vegetable oils using nickel as catalyst gives edible fats (margarine and vanaspati ghee)

(ii) Hydroformylation of olefins yields aldehydes which further undergo reduction to give alcohols.

H2 + CO + RCH = CH2→ RCH2CH2CHO

H2 + RCH2CH2CHO→RCH2CH2CH2OH

Ortho and parahydrogen

  • Normal hydrogen consists of a mixture of two forms : Ortho hydrogen (75%) and para hydrogen (25%). In ortho hydrogen, the two H atoms constituting H2 molecule have same nuclear spin, but in para hydrogen the nuclear spin of the two H atoms of H2 molecule is
  • The two forms, ortho and para hydrogen possess same chemical properties but different physical properties.
  • Thermal conductivity of para hydrogen is 50% more than ortho hydrogen.
  • Pure para hydrogen can be obtained at low temperature but pure ortho hydrogen cannot be obtained.

Atomic hydrogen

  • It is formed by dissociation of hydrogen molecules. The dissociation is caused by passing molecular hydrogen through the electrical arc struck between tungsten electrodes.

H2(g) 2H(g); DH = 436 kJ

  • It is highly energetic, highly unstable and has a very short life span of about 0.3 sec. It is highly reactive and cannot be stored as such.
  • It is used in atomic hydrogen torch.

Nascent hydrogen

  • It is the form of hydrogen at the moment of its generation from chemical reaction in aqueous solutions. It is represented generally by [H].
  • In general, the reactivity order is : H2 < Nascent hydrogen < Atomic hydrogen.

HYDRIDES

Ionic or Saline Hydrides

  • Alkali metals and alkaline earth metals combine with dihydrogen to form Ionic or Saline Hydrides.
  • LiH, BeH2 and MgH2 have significant covalent character due to very high polarizing power of Li+, Be+2 & Mg+2.
  • BeH2 and MgH2 are polymeric in structure.
  • The ionic hydrides are crystalline, non-volatile and nonconducting in solid state. However, their melts conduct electricity.
  • On electrolysis dihydrogen is liberated at anode

Hydrolysis of Ionic or Saline hydrides produces dihydrogen gas and the reaction is highly exothermic in nature

NaH(s) + H2O(aq)→NaOH(aq) + H2(g)

  • Lithium hydride is rather unreactive at moderate temperatures with O2 or Cl2. It is, therefore, used in the synthesis of other useful hydrides, e.g.,

8LiH + Al2Cl6→2LiAlH4 + 6LiCl          ;           2LiH + B2H6→2LiBH4

Covalent or Molecular Hydride

  • Dihydrogen reacts with most of the p-block elements to form covalent or molecular hydrides.
  • Being covalent, they are volatile in nature.

Molecular hydrides are further classified according to the relative numbers of electrons and bonds in their Lewis structure into :

(i) Electron-deficient (ii) electron-precise, and (iii) electron-rich hydrides.

Electron deficient hydrides :  In general, Group 13 elements combine with dihydrogen to form electron deficient hydrides. Due to incomplete octet and presence of vacant p-orbital they can accept electrons hence they act as Lewis acids. e.g. Diborane (B2H6)

Electron-precise hydrides :In general, Group 14 elements combine with dihydrogen to form electron precise hydrides. They have complete octet. which are tetrahedral in geometry. e.g., CH4

Electron-rich hydrides : In general, Group 15-17 elements combine with dihydrogen to form electron rich hydrides. Due to the presence of lone pair they can donate electrons and behave as good Lewis bases e.g. NH3 (1 lone pair), H2O (2 lone pair) & HF (3 lone pair).

The presence of lone pairs on highly electronegative atoms like N, O and F in hydrides results in hydrogen bond formation between the molecules. This leads to the association of molecules.

Metallic or Non-stoichiometric (or Interstitial) Hydrides

  • In general, d-block and f-block elements combine with dihydrogen to form metallic, non-stoichiometric (or interstitial) hydrides.
  • The metals of group 7, 8 and 9 do not form hydride
  • from group 6, only chromium forms CrH
  • They show thermal and electrical conductivity but not as saline or ionic hydrides
  • They are nonstoichiometric, being deficient in hydrogen. For example, LaH87, YbH2.55, TiH1.5–1.8, ZrH1.3–1.75, VH0.56, NiH0.6–0.7, PdH0.6–0.8 etc.

(In such hydrides, the law of constant composition does not hold good.)

  • recent studies have shown that except for hydrides of Ni, Pd, Ce and Ac, other hydrides of this class have lattice different from that of the parent metal.
  • Some of the metals (e.g., Pd, Pt) can accommodate a very large volume of hydrogen and, therefore, can be used as its storage media. This property has high potential for hydrogen storage and as a source of energy.

WATER

Human body has about 65% water and some plants have as much as 95% water.

Estimated World Water Supply

Rivers < Atmospheric water vapour < Soil Moisture < Saline lake and inland seas < Lakes < Ground Water < Polar ice and glaciers < Oceans

Physical Properties of Water

  • It is a colour less and tasteless liquid.
  • Due to strong intermolecular H-bonding, water (H­2O) has high freezing point, high boiling point, high heat of vaporisation and high heat of fusion in comparison to H2S and H2
  • Boiling Point, Melting Point, Enthalpy of formation, Enthalpy of Vaporisation, Enthalpy of fusion, Density & Viscosity of H2O < D2O
  • Dielectric Constant of H2O > D2O
  • It is an excellent solvent for transportation of ions and molecules required for plant and animal metabolism.
  • Due to hydrogen bonding with polar molecules, even covalent compounds like alcohol and carbohydrates dissolve in water.

Structure of Water

In gas and liquid phase
  • In the gas phase water is a bent molecule with a bond angle of 104.5°, and O–H bond length of 95.7 pm.
  • Due to angular shape, it is a polar molecule with dipole moment 1.83 D, and dielectric constant (82.5).
  • Water is made of hydrogen and oxygen elements. The composition of hydrogen and oxygen by mass is 1 : 8 (determined by Morley and Dumas) and that by volume is 2 : 1 (determined by Hoffmann).
In Solid Phase
  • The crystalline form of water is ice.
  • At atmospheric pressure ice crystallises in the hexagonal form,
  • At very low temperatures it condenses to cubic form.
  • Density of ice is less than that of water. Therefore, an ice cube floats on water.
  • In winter season ice formed on the surface of a lake provides thermal insulation which ensures the survival of the aquatic life. This fact is of great ecological significance.
Structure of Ice

Ice has a highly ordered three dimensional hydrogen bonded structure.

  • X-rays shows that each oxygen atom is surrounded tetrahedrally by four other oxygen atoms at a distance of 276 pm.
  • Due to strong intermolecular H-bonding in ice it forms cage like structure which can accommodate small size atoms in the interstitial space.

Chemical Properties of Water

Amphoteric Nature:

(a). As a Bronsted Acid

H2O(l) + NH3(aq) OH(aq) + NH4+(aq)

(b). As a Bronsted Base

H2O(l) + H2S(aq) H3O+(aq) + HS(aq)

(c). Both as Bronsted Acid as well as Bronsted Base (Auto-protolysis or self-ionization)

Redox Reactions Involving Water:

(a). Water acts as an oxidising agent with highly electropositive metals

2H2O(l) + 2 Na(s)→2NaOH(aq) + H2 (g)

(b). Water acts as a reducing agent during photosynthesis

6CO2(g) + 12H2O(l)→C6H12O6(aq) + 6H2O(l) + 6O2(g)

(c) . Water acts as a reducing agent with F2

2F2(g) + 2H2O(l)→4H+(aq) + 4F(aq) + O2(g)

Hydrolysis Reaction :

P4O10(s) + 6H2O(l)→4H3PO4(aq)

SiCl4(l) + 2H2O(l)→SiO2(s) + 4HCl(aq)

N3–(s) + 3H2O(l)→NH3(g) + 3OH(aq)

Hydrates Formation:

(i) Coordinated water                 e.g.,

(ii) Interstitial water                   e.g.,      BaCl2.2H O

(iii) Hydrogen-bonded water      e.g.,

Hard and Soft Water

Hard Water : Presence of calcium and magnesium salts in the form of hydrogen carbonate, chloride and sulphate in water makes water ‘hard’. Hard water does not give lather with soap.

Soft Water : Water free from soluble salts of calcium and magnesium is called Soft water. It gives lather with soap easily.

Hard water forms scum/precipitate with soap. Soap containing sodium stearate (C17H35COONa) reacts with hard water to precipitate out Ca/Mg stearate.

2C17H35COONa(aq) + M2+(aq)→ (C17H35COO)2 M↓+2Na+ (aq) ; M is Ca/Mg

The hardness of water is of two types:

(i) Temporary hardness                            (ii) Permanent hardness.

Temporary Hardness

Temporary hardness is due to the presence of magnesium and calcium hydrogen carbonates. It can be removed by :

(i) Boiling:.

(ii) Clark’s method: Removal of hydrogen carbonates of calcium and magnesium using Ca(OH) 2  or lime in the form of  calcium carbonate and magnesium hydroxide is called as Clark’s method.

Ca(HCO3)2 + Ca(OH)2→2CaCO3 ↓+2H2O

Mg(HCO3)2 + 2Ca(OH)2→2CaCO3↓+Mg(OH)2↓+2H2O

Permanent Hardness

It is due to the presence of soluble salts of magnesium and calcium in the form of chlorides and sulphates in water. Permanent hardness is not removed by boiling. It can be removed by the following methods:

(i)   Treatment with washing soda (sodium carbonate): Washing soda reacts with soluble calcium and magnesium chlorides and sulphates in hard water to form insoluble carbonates.

MCl2 + Na2CO3→MCO3↓ + 2NaCl (M = Mg, Ca)

MSO4 +Na2CO3→MCO3↓+ Na2SO4

(ii) Calgon’s method: Sodium hexametaphosphate (Na6P6O18), commercially called ‘calgon’, when added to hard water, the following reactions take place.

The complex anion keeps the Mg2+ and Ca2+ ions in solution.

(iii) Ion-exchange method: This method is also called zeolite/permutit process. Hydrated sodium aluminium silicate is zeolite/permutit. For the sake of simplicity, sodium aluminium silicate (NaAlSiO4) can be written as NaZ. When this is added in hard water, exchange reactions take place.

2NaZ(s) + M2+(aq)→MZ2(s) + 2Na+(aq) (M= Mg, Ca)

Permutit/zeolite is said to be exhausted when all the sodium in it is used up. It is regenerated for further use by treating with an aqueous sodium chloride solution.

MZ2 (s) +2NaCl(aq)→2NaZ(s) + MC2(aq)

(iv) Synthetic resins method: Nowadays hard water is softened by using synthetic cation exchangers. This method is more efficient than zeolite process. Cation exchange resins contain large organic molecule with – SO3H group and are water insoluble. Ion exchange resin (RSO3H) is changed to RNa by treating it with NaCl. The resin exchanges Na+ ions with Ca2+ and Mg2+ ions present in hard water to make the water soft. Here R is resin anion.

2RNa(s) + M2+(aq)→R2M(s) + 2Na+ (aq)

The resin can be regenerated by adding aqueous NaCl solution.

Pure de-mineralised (de-ionized) water free from all soluble mineral salts is obtained by passing water successively through a cation exchange (in the H+ form) and an anion exchange (in the OHform) resins:

2RH(s) + M2+ (aq) MR2(s) + 2H+(aq)

In this cation exchange process, H+ exchanges for Na+, Ca2+, Mg2+ and other cations present in water. This process results in proton release and thus makes the water acidic. In the anion exchange process OH exchanges for anions like  etc. present in water.

RNH2 (s) + H2O(l) . OH(s)

.OH(s) + X(aq) .X(s) + OH(aq)

OHions, thus, liberated neutralise the H+ ions set free in the cation exchange.

H+(aq) + OH(aq)→H2O(l)

HYDROGEN PEROXIDE (H2O2)

Preparation

It can be prepared by the following methods.

(i)   Acidification of Barium Peroxide: Barium peroxide is acidified using dilute sulphuric acid and excess water is removed by evaporation under reduced pressure to gives hydrogen peroxide.

BaO2.8H2O(s) + H2SO4(aq)→BaSO4 (s) + H2O2(aq) + 8H2O(l)

(ii) Electrolytic Oxidation : Electrolytic oxidation of acidified sulphate solutions at high current density forms Peroxodisulphate which on hydrolysis yields hydrogen peroxide.

This method is now used for the laboratory preparation of D2O2.

K2S2O8(s) + 2D2O(l)→2KDSO4(aq) + D2O2(l)

(iii)  Industrial Preparation : Industrially it is prepared by the autooxidation of 2-alklylanthraquinols.

In this case 1% H2O2 is formed. It is extracted with water and concentrated to ~30% (by mass) by distillation under reduced pressure. It can be further concentrated to ~85% by careful distillation under low pressure. The remaining water can be frozen out to obtain pure H2O2.

Physical Properties

In the pure state H2O2 is an almost colourless (very pale blue) liquid.

H2O2 is miscible with water in all proportions and forms a hydrate H2O2.H2O (mp 221K). A 30% solution of H2O2 is marketed as ‘100 volume’ hydrogen peroxide. It means that one millilitre of 30% H2O2 solution will give 100 mL of oxygen at STP. Commercially marketed sample is 10 V, which means that the sample contains 3% H2O2.

Structure

Hydrogen peroxide has a non-planar structure. The molecular dimensions in the gas phase and solid phase are

  (a) Gas phase (dihedral angle is 111.5°)                 (b) Solid phase (dihedral angle is 90.2°)

Chemical Properties

It acts as an oxidising as well as reducing agent in both acidic and alkaline media. Simple reactions are described below.

(i)   Oxidising action in acidic medium

2Fe2+(aq) + 2H+(aq) + H2O2(aq)→2Fe3+(aq) + 2H2O(l)

PbS(s) + 4H2O2(aq)→PbSO4(s) + 4H2O(l)

(ii) Reducing action in acidic medium

(iii) Oxidising action in basic medium

2Fe2+ + H2O2→2Fe3+ + 2OH

Mn2+ + H2O2 → Mn4+ + 2OH

(iv) Reducing action in basic medium

I2 + H2O2 + 2OH→2I + 2H2O + O2

Storage

H2O2 decomposes slowly on exposure to light.

2H2O2(l)→2H2O(l) + O2(g)

In the presence of metal surfaces or traces of alkali (present in glass containers), the above reaction is catalysed. It is, therefore, stored in wax-lined glass or plastic vessels in dark. Urea can be added as a stabiliser. It is kept away from dust because dust can induce explosive decomposition of the compound.

HEAVY WATER, D2O

  • It is extensively used as a moderator in nuclear reactors and in exchange reactions for the study of reaction mechanisms.
  • It can be prepared by exhaustive electrolysis of water or as a by-product in some fertilizer industries.

It is used for the preparation of other deuterium compounds, for example:

CaC2 + 2D2O→C2D2 + Ca(OD)2

SO3 + D2O→D2SO4

Al4C3 + 12D2O→3CD4 + 4Al(OD)3

Hydrogen Economy : The basic principle of hydrogen economy is the transportation and storage of energy in the form of liquid or gaseous dihydrogen. Advantage of hydrogen economy is that energy is transmitted in the form of dihydrogen and not as electric power.

Initially 5% dihydrogen has been mixed in CNG for use in four-wheeler vehicles.