UNIT 9 HYDROGEN
Syllabus
· Position of hydrogen in the periodic table.
· Occurrence
· Isotopes
· Preparation
· Properties and uses of hydrogen
· Hydrides – ionic, covalent and interstitial
· Physical and chemical properties of water
· Heavy water
· Hydrogen peroxide –Preparation, reactions and structure
· Hydrogen as a fuel
Hydrogen is the most abundant element in the universe. It was detected by Paracelus during the reaction of sulphuric acid on iron in the early sixteenth century. Later on in 1766 Henry Cavendish prepared it in the pure form and studied many of its properties. The name hydrogen meaning the substance capable of forming water was given by Lavoisier in 1783.
Hydrogen is the lightest element and is the first member of the periodic table. The chemistry of hydrogen is quite interesting as it shows unique properties due to unique electronic configuration, 1s1.
UNIQUE POSITION OF HYDROGEN IN THE PERIODIC TABLE
Hydrogen is unique among the elements in several ways. It is lightest, simplest and the first element of the periodic table. It is the most abundant element in the universe. In its properties, it behaves like alkali metals as well as halogens. Hydrogen is grouped along with alkali metals and is placed as the first element of the first Group in the periodic table.
Resemblance with alkali metals
1. Electronic configuration : Hydrogen like alkali metals has just one electron in its valency shell.
2. Electropositive character : The loss of one electron from hydrogen yields H+ ions just like alkali metals.
H(g) ® H+ + e-
M(g) ® M+(g) + e-
3. Valency : Like the alkali metals, hydrogen is monovalent and possesses strong reducing properties, typical of electropositive elements.
4. Liberation at the cathode : During electrolysis of water hydrogen is liberated at the cathode. Similarly, during electrolysis of molten alkali metal salts , alkali metal is deposited at the cathode.
5. Reducing Nature : Hydrogen removes oxygen from oxides and acts as strong reducing agents like alkali metals.
Fe3O4 + 4 H2 ® 3 Fe + 4 H2O
B2O3 + 6 K ® 2 B + 3 K2O
6. Affinity for electronegative Elements : Hydrogen like alkali metals has strong affinity for electronegative elements
such as oxygen, sulphur, halogens etc. and combine with them to form respective compounds. E.g., Na2O, Na2S, NaCl (compounds of alkali metals) H2O, H2S, HCl (compounds of hydrogen).
Although hydrogen behaves like alkali metals in some respects, it differs from alkali metals in the following respects:
Differences of Hydrogen from Alkali metals
1. Ionic size : The size of H+ ion is very small as compared to the size of alkali metal ions.
2. Ionisation energy : Ionisation energy of hydrogen (1312 kJ/mol) as compared to those of alkali metals(520-376 kJ/mol) is high. Hence compounds of hydrogen are covalent with polar nature.
3. Anion formation : Hydrogen can form anions, while alkali metals cannot.
4. Atomicity : Hydrogen is diatomic whereas alkali metals are monoatomic.
5. Non-metallic character : Hydrogen is a typical non-metal while alkali metals are metallic in nature.
Resemblance with Halogens
1. Electronic configuration : Like halogens, hydrogen also requires one more electron to attain stable configuration of the nearest noble gas.
2. Ionisation Energy : The ionisation energy of hydrogen is of the same order as that of halogens ( Cl = 1252 kJ/mol).
3. Liberation at the anode : Like halogens , hydrogen is also liberated at the anode when molten salts with alkali metal(M) are electrolysed.
4. Atomicity : In the elementary state, hydrogen and halogen are diatomic molecules.
5. Non-metallic character : Like halogens, hydrogen is typical non-metallic element.
6. Electronegative character : Hydrogen accepts an electron to form monovalent anions like halogens.
H + e- ® H-
Cl + e- ® Cl-
Although hydrogen behaves like halogens in some respects, it differs from halogens in the following respects.
Differences from Halogens
1. Reactivity : Hydrogen is less reactive while halogens are highly reactive at room temperature.
2. Tendency to form negative ion : Tendency of hydrogen to form negative ion is very low as compared to halogens.
3. Hydrogen bonding : Hydrogen is capable of forming hydrogen bonds while halogens do not show any such properties.
Conclusion
From the foregoing discussions, it is evident that hydrogen is unique element in the sense it resembles both alkali metals and halogens. At the same time hydrogen differs both the families in certain properties. However, it has been placed at the top of the alkali metal family. This is because of the resemblance in its electronic configuration (1s1) with alkali metals.
DIHYDROGEN, H2
Occurrence
Hydrogen is the most abundant element in the universe(70% of the universe’s total mass) . The giant planets Jupiter and Saturn consists mostly of hydrogen. About half the mass of the sun and some other stars is made up of hydrogen. It is the third most abundant element on the surface of the globe. In the combined form it constitutes 15.4% of the earth’s crust and the oceans. Earth does not possess enough gravitational pull to retain light H2 molecules , so it is not in our atmosphere.
In the crustal rocks it is 10th in the order of abundance ~ 0.15% by weight.
ISOTOPES OF HYDROGEN
The three isotopes of hydrogen are protium (11H), deuterium (21H or D) and tritium (31H or T).
The predominant form is protium whose atoms are made up of a single proton and a single electron. The deuterium nucleus is made up of a proton and a neutron, giving it a mass number of 2. Tetrestrial hydrogen contains 0.0156% of deuterium most in the form of HD.
The tritium (31H ) concentration is about 1 atom per 1018 atoms of protium. It is radioactive and emits low energy b-particles ( t ½ 12.33 y).
The characteristic atomic properties of the three isotopes are given in thr TABLE.
TABLE 1
Atomic properties of Isotopes of Hydrogen
Property | H | D | T |
Relative atomic mass | 1.007825 | 2.014102 | 3.016049 |
Nuclear spin quantum number | ½ | 1 | ½ |
Radioactive stability | Stable | Stable | b, t½ = 12.33 y |
The atomic properties of the three isotopes indicate that they can be used in isotope studies, radioactive tracer studies and NMR spectroscopy.
PREPARATION AND COMMERCIAL PRODUCTION OF HYDROGEN
Small scale Production of Dihydrogen
1. In the laboratory , dihydrogen can be prepared very conveniently by the action of dilute sulphuric acid on granulated zinc even at room temperature.
Zn(s) + H2SO4 (dil) ® ZnSO4 + H2 (g)
The apparatus used for this purpose is shown in Fig.
Laboratory preparation of dihydrogen
Granulated zinc is taken in a Woulf’s bottle fitted with a thistle funnel in one mouth and a delivery tube in the other. Dilute sulphuric acid is poured through the funnel. The reaction takes at room temperature and no heating is required. The evolved hydrogen gas is collected by the downward displacement of water.
Following points are to be remembered in the above method of preparation of hydrogen.
i. Pure zinc is not used in the preparation of dihydrogen because it is non-porous in nature and therefore, the reaction between Zn and H2SO4 is slow. The presence of impurities makes zinc porous, which helps in constituting electrochemical couple and speed up the reaction.
ii. Con. Sulphuric acid is not used because it acts as acid as well as oxidising agent and consequently get reduced during the process to liberate SO2.
2. It can also be prepared by the action of zinc with aqueous alkali.
Zn + 2 NaOH ® Na2ZnO2 + H2
Sodium zincate
COMMERCIAL PRODUCTION OF DIHYDROGEN
1. By the Electrolysis of Acidified water
Hydrogen is manufactured by the electrolysis of acidified water using platinum electrodes.
Hydrogen of high purity (> 99.95% ) is obtained by electrolysing warm aqueous barium hydroxide between nickel electrodes.
2. From hydrocarbons and coke
Hydrogen is obtained by the reaction of steam on hydrocarbons or coke at high temperatures.
The mixture of CO and H2 is used in the synthesis of methanol and a number of hydrocarbons, it is also called synthesis gas or ‘syngas’. Nowadays ‘syngas’ is produced from sewage , saw dust , scrap wood , newspapers etc. The process of producing ‘syngas’ from coal is called ‘coal gasification’.
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.
3. Relatively smaller quantities of dihydrogen are obtained by passing a 1 : 1 molar mixture of vapourised methanol and water over a ‘base metal chromite’ type catalyst at 673 K. The mixture of hydrogen and carbon monoxide obtained is made to react with steam to give CO2 and more hydrogen.
The resulting gas mixture is compressed to about 25 atm. pressure and is passed into water. CO2 dissolves in water under pressure, while dihydrogen is released and is collected.
4. Hydrogen can also produced as a by-product of the brine electrolysis for the manufacture of chlorine and sodium hydroxide. During electrolysis , the reactions that take place are :
at anode : 2 Cl- (aq) ® Cl2 + 2 e-
at cathode : 2 H2O(ℓ) + 2 e- ® H2(g) + 2 OH- (aq)
The overall reaction is :
2 Na+(aq) + 2 Cl- (aq) + 2 H2O(ℓ) ®
Cl2(g) + H2(g) + 2 Na+(aq) + 2 OH- (aq)
PROPERTIES OF DIHYDROGEN
Physical properties
Dihydrogen is a colourless , odourless and tasteless gas. The physical parameters with those for dideuterium are given in TABLE 2. The vaues for for deuterium are substantially higher as compared to diprotium. The relatively inert behaviour of dihydrogen at room temperature is due to high enthalpy of H- H bond and being the highest for a single bond between any two elements. Significant decomposition of H2 , into its atoms occurs only above 2000 K (0.081%) which increases to 95.5% at 5000 K.
Dihydrogen has two nuclear spin isomers called ortho and para–dihydrogen. Their nuclei possess parallel and antiparallel spins respectively.
The two isomers have 1 and 0 nuclear spin respectively. Conversion of one isomer into the other is a slow process. Para- hydrogen with lower energy is favoured at low temperatures. Above 0 K , the equilibrium concentration of the ortho isomer increases until a limiting 3 : 1 proportion of ortho : para is established. It is possible to obtain pure para-hydrogen . It is never possible to obtain a sample containing more than 75% ortho hydrogen.
The physical properties of dihydrogen are not significantly affected by nuclear spin isomerism. Some of the notable difference are :
(i) The thermal conductivity of para-hydrogen is 50% greater than that of the ortho-hydrogen.
(ii) The melting point of para hydrogen is 0.15 K below that of hydrogen containg 75% ortho-hydrogen
CHEMICAL PROPERTIES
1. Reaction with halogens
Dihydrogen combines with halogens (X) to give hydrogen halides. The reactivity of halogens towards dihydrogen decreases in the order F2 > Cl2 > Br2 > I2. The reaction with fluorine takes place even in dark ; with iodine a catalyst is required.
2. Reaction with oxygen
Dihydrogen combines with dioxygen either at 970 K or upon passing electric discharge to form water vapours. The reaction is highly exothermic
3. Reaction with nitrogen
Dihydrogen and dinitrogen combine to form ammonia under following conditions.
The method is used to manufacture ammonia by Haber process.
4. Action with metals
Dihydrogen combines directly with alkali metals, alkaline earth metals and certain rare earh metals to form corresponding compounds known as hydrides. In these hydrides hydrogen exists in - 1 state. Due to their ionic nature, these hydrides are also referred to as salt like hydrides.
The metals like iron , nickel, chromium, palladium etc. absorb dihydrogen in almost stoichiometric quantities. This phenomenon is known as occlusion and the compounds thus formed are known as interstitial hydrides or metallic hydrides. During the formation of these compounds hydrogen atoms occupy the interstitial sites present in the crystals of metals.
5. Reaction with metal ions : Dihydrogen reduces some metal ions in aqueous solutions into corresponding metals. For example,
Pd2+(aq) + H2(g) ® Pd(s) + 2 H+(aq)
6. Reaction with metal oxides (reducing property)
When dihydrogen gas is passed over heated metal oxides it reduces them into corresponding metals. Thus it acts as a strong reducing agent. Some examples showing the reducing action of dihydrogen are as follows.
7. Reaction with carbon monoxide
Dihydrogen reacts with carbon monoxide at 700 K and 200 atmosphere pressure in presence of ZnO/ Cr2O3 catalyst to form methanol.
8. Reaction with unsaturated hydrocarbons
In presence of finely divided nickel (or platinum) dihydrogen reacts with unsaturated hydrocarbons such as ethylene, acetylene etc. at 473 K to form saturated hydrocarbons.
This reaction (hydrogenation) is used for the conversion of polyunsaturated oils into edible fats. Polyunsaturated vegitable oils have too many C=C bonds which undergo oxidation and the oil becomes unpleasant in taste, that is , it becomes rancid. Hydrogenation reduces the number of double bonds , but does not completely eliminate them. Under controlled conditions, edibile oils and margarine can be prepared by hydrogenation of vegetable oils like soyabean, corn, cotton seed oils. Nickel is used as the catalyst. Vanaspathi – an edible cooking medium is prepared on large scale by the hydrogenation process.
USES OF DIHYDROGEN
1. The most important use of dihydrogen is in the manufacture of ammonia by Haber process.
2. Dihydrogen is used in the manufacture of methanol.
3. Direct reaction of dihydrogen with chlorine is used for the preparation of hydrogen chloride.
4. Dihydrogen is used in the manufacture of metal hydrides.
5. In metallurgical processes , it is used to reduce heavy metal oxides to metals.
6. An important use of hydrogen is in the atomic hydrogen and oxy-hydrogen torches for cutting and welding. Dihydrogen is dissociated with the help of an electric arc and the hydrogen atoms obtained are allowed to recombine on the surface to be welded. High temperature of 4000 K is generated.
7. Hydrogen is used as a fuel in space rockets and guided missiles.
8. Hydrogen is used in H2-O2 fuel cells
Problem
01. Comment on the reactions of dihydrogen with (i) Chlorine
(ii) sodium (iii) and copper (II) oxide.
HYDRIDES
Dihydrogen , under certain reaction conditions , combines with almost all elements , except noble gas to form binary compounds, called hydrides. If ‘ E ’ is the symbol of an element then hydride can be expressed as EHx ( e.g., MgH2) or EmHn ( e.g., B2H6).
The hydrides are classified into threew categories :
i) Ionic or saline or salt like hydrides.
ii) Covalent or molecular hydrides
iii) Metallic or non-stoichiometric hydrides
1. Saline hydrides (Ionic hydrides)
These are compounds of hydrogen with most of the s-block elements. They are non-volatile, non-conducting , crystalline solids. They are also referred to as ‘salt like’ hydrides. However, BeH2 and MgH2 have covalent polymeric structures.
The binary hydrides of alkali metals ( LiH, NaH, KH , RbH, CsH) have rock salt structures. The thermal stability of alkali metal hydrides decreases from LiH to CsH. The order of stability in the alkaline earth metal hydrides is :
CaH2 > SrH2 > BaH2
Electrolysis of solutions of saline hydrides in molten alkali halides produces hydrogen at the anode which confirms the existence of the hydride ion H- ion.
Saline hydrides react explosively with water
The fires so produced cannot be extinguished by carbon dioxide as it gets reduced by hot metal hydride. Only sand is useful as it is a solid.
Alkali metal hydrides are used as reagents for preparing other hydride compounds, e.g. LiAlH4 and LiBH4.
While the hydrides of heavier members of each group are supposed to have ionic bonds ; there seems to be significant covalent character in the lighter metal hydrides ( LiH, MgH2 and BeH2)
2. Molecular hydrides( Covalent hydrides)
Dihydrogen forms molecular compounds with most of p-block elements. Most of the familiar examples are CH4 , NH3 , H2O and HF. For convenience hydrogen compounds of non-metals have also been considered as hydrides. Being covalent, they are volatile compounds.
The systematic name for molecular hydrides are usually formed from the name of the element and the suffix –ane . Phosphane for PH3, Oxidane for H2O and azane for NH3.
Molecular hydrides are further classified according to the relative numbers of electrons and bonds in their Lewis structures.
i) electron deficient molecular hydride.
ii) Electron precise molecular hydride
iii) Electron rich molecular hydride.
An electron deficient molecular hydride has too few electrons for writing its conventional Lewis structure. Diborane B2H6 is an example.
Electron-precise compounds are formed by elements of group 14. The molecules are tetrahedral, for example CH4. The bond distance increases on going down the group.
Ammonia and water are electron-rich hydrides. The excess electrons being present as lone pairs of electrons ( ammonia – one lone pair ,water- 2 lone pairs and hydrogen fluoride – 3 lone pairs).
The presence of lone pairs and highly electronegative atoms like nitrogen, oxygen and fluorine in electron-rich hydrides result in hydrogen bond formation between the molecules. This leads to association of molecules.
3. Metallic Hydride or Non-stiochiometric (Interstitial ) Hydrides
These hydrides are formed by d-block (Groups 3, 4 and 5) and f-block elements. In Group 6 chromium alone forms the hydride CrH. The metals of group 7, 8 and 9 do not form hydrides. The region of the periodic table from group 7 to 9 is referred to as the Hydride Gap. These hydrides conduct heat and electricity though not as efficiently as their parent metals do. Unlike saline hydrides, they are almost always non-stoichiometric , being deficient in hydrogen. For example, LaH2.87 , YbH2.55 , TiH1.5 –1.8 , ZrH1.3 – 1.75, VH0.56 , NiH0.6–0.7 , PdH0.6–0.8 etc. In such hydrides , law of constant composition does not hold good.
Earlier it was thought that in these hydrides , hydrogen occupies interstices in metal lattice producing distortion without any change its type. Consequently, they were termed as interstitial hydrides. However, 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. The property of absorption of hydrogen on transition metals is widely used in catalytic reduction/hydrogenation reactions for the preparation of large number of compounds. Some of the metals (e.g., Pd , Pt ) can accommodate a very large volume of hydrogen and therefore , can be used as storage media. This property has high potential for hydrogen storage and as source of energy.
Problem
02. Would you expect the hydrides of N , O and F to have lower boiling points than the hydrides of their subsequent group members ? Give reasons.
03. Can phosphorus with outer electronic configuration 3s23p3 form PH5 ?
WATER
Water is the most common, abundant and easily obtainable of all chemical compounds. It is crucial compound for man’s survival. It can be easily transformed from liquid to solid and to gaseous states.
Physical properties
It is a colourless and volatile liquid. It has many unusual properties. Some of the physical properties of H2O and heavy water D2O are listed below:
Property | H2O | D2O |
Molecular mass (g/mol) | 18.0151 | 20.0276 |
Melting point (K) | 273.0 | 276.8 |
Boiling point (K) | 373.0 | 374.4 |
Enthalpy of formation kJ/mol | - 285.9 | - 294.6 |
Enthalpy of vaporisation(373 K) kJ/mol | 40.66 | 41.61 |
Enthalpy of fusion kJ/mol | 6.01 | - |
Density (g/cm3) | 1.0000 | 1.1059 |
Temp.of maximum density | 276.98 k | 284.61 k |
Viscosity (centipoise) | 0.893 | 1.107 |
Dielectric constant | 78.39 | 78.06 |
Electrical conductivity (293 K mho/cm) | 5.7 x 10-8 |
The peculiar properties of water in condensed phases (liquid and solid state) are due to the presence of existence of hydrogen bonding between water molecules. The abnormally high freezing point, boiling point , heat of vaporisation and heat of fusion (compared to the hydrides of other elements of the same group of the periodic table, e.g. H2S , H2Se ) are due to hydrogen bonding.
Water has a higher specific heat , thermal conductivity and surface tension than most other liquids. These properties allow water to play a vital role in biosphere. The high heat of vaporisation and high heat capacity of water are responsible for moderation of the climate and body temperature of living organisms. Water is an excellent solvent for transportation of ions and molecules needed for plant and animal metabolism.
Even covalent organic compounds like alcohols and carbohydrates have high solubility in water because of their ability to form hydrogen bonds with water. Under very high pressure and temperature water behaves as non-polar solvent and dissolves organic compounds while common inorganic salts are not dissolved.
When water freezes it forms ice- the crystalline form of water. Nine structurally different forms of ice are known. At atmospheric pressure ice crystallizes in the normal hexagonal form, Ih ,but at very low temperatures it condenses to cubic form. Ice (Ih) has lesser density than liquid phase with which it is in equilibrium. This property is of crucial significance for the preservation of aquatic life.
Many salts crystallize from aqueous solutions as hydrated salts. Such association of water is essentially of five types : co-ordinated water, hydrogen bonded water , lattice water, zeolite water and clathrate water.
STRUCTURE OF WATER
The formula of a single molecule of water is H2O. The molecule consists of two hydrogen atoms joined to an oxygen atom by covalent bonds (Fig a).
Structure of Water molecule
The water molecule assumes a tetrahedral shape as suggested by VSEPR theory. The two bond pairs and the two lone pairs present in the central oxygen atom orient themselves in a tetrahedral geometry as shown in Fig(b). Consequently water molecule has an angular or bent structure. The repulsive interaction of the lone pairs and bond pairs decreases the bond angle from 109°28’ to 104.5°. The two OH bonds are equal in length and the O - H bond length is 95.7 pm. Oxygen is highly electronegative and has strong tendency to pull the shared pairs of electrons towards its own side. This results in the development of partial negative charge on oxygen and partial positive charge on the two hydrogen atoms as shown in Fig c. The bond moments of O - H bonds causes the molecule to behave as a permanent dipole. The dipole moment of H2O molecule is found to be 1.85 D which confirms its polar nature.
In the liquid state , water molecules are held together by intermolecular hydrogen bonding as shown in Fig.
Hydrogen bonding in liquid water
Each oxygen can form two hydrogen bonds with the hydrogen atoms of the neighbouring molecules. Thus in liquid state, water consists of aggregates of varying number of H2O molecules held together by hydrogen bonds. Some free H2O molecules also exist. In fact there is a dynamic equilibrium between the aggregates of molecules and free molecules. The free water molecules continuously form aggregates which continuously collapse and change into free molecules. The presence of aggregates of molecules due to hydrogen bonding affects several properties of water to a considerable extent. The existence of water in liquid state , its exceptionally high boiling point as compared to other hydrides of the elements of Group 16 and many more curious properties of water may be attributed to the presence of hydrogen bonding in it.
In solid state , the water molecules come still closer and get arranged tetrahedrally. The solid form of water, i.e., exists in different crystalline forms depending upon the conditions used for freezing water. The structure of normal hexagonal is shown in Fig. In this type of ice, each oxygen atom is tetrahedrally surrounded by four other oxygen atoms through a hydrogen atom. Each hydrogen is covalently bonded to one oxygen atom and linked to another oxygen atom by a hydrogen bond. This type of packing leads to an open cage structure with large open spaces. This is why the density of ice is less than that of water and it floats on water.
Structure of normal hexagonal ice
Each oxygen atom is bonded to four hydrogen atoms. The covalent bonds are represented by shorter full lines, and the weaker hydrogen bonds by longer broken lines between O and H atoms.
This fact is of great ecological significance. The ice layer , formed on the surface of a lake , for example, in winter , does not sink to the bottom. It provides thermal insulation for water below it and this ensures the survival of aquatic life.
Chemical properties of water
Water shows a varied and unique chemical behaviour.
1. Amphoteric behaviour :It has the ability to act as an acid or a base, in Bronsted sense, that is, it behaves as an amphoteric substance. It can for example, act as an acid towards ammonia :
H2O(ℓ) + NH3 (aq) ® OH-(aq) + NH4+(aq)
And as a base towards H2S :
H2S(aq) + H2O(ℓ) ® H3O+(aq) + HS-(aq)
In general, water can act as a base towards acids stronger than itself and as an acid towards a base stronger than it. The auto-protolysis of water represented by :
H2O(ℓ) + H2O(ℓ) ⇌ H3O+(aq) + OH-(aq)
Acid 1 base 2 acid 2 base 1
is of great importance in acid-base chemistry.
2. Redox Reactions Involving Water : Another important class of reactions in aqueous chemistry are oxidation – reduction reactions. Water is reduced to H2 by metals with E° value of the redox couple Mn+/M below - 0.41 V .
2 H2O(l) + 2 e ® 2 OH- (aq) + H2(g) : E° = - 0.41 V
for [OH-] = 10-7 M Water can also be oxidised as shown by the equation :
O2(g) + 4 H+(aq) + 4 e- ® 2 H2O : E° = + 0.82 V
or [H+] = 10-7 M
Some typical examples of this type of reaction are oxidation of water to oxygen during photosynthesis and oxidation by fluorine.
2 F2(g) + 2 H2O(l) ® 4 H+ (aq) + 4 F- (aq) + O2(g)
3. Hydrolysis Reaction : Water with high dielectric constant has a very strong solvating character. It acts as an excellent solvent. Because of this great affinity for many elements for oxygen , hydrolysis of ionic or covalent compounds occurs in water, e.g.,
P4O10(s) + 6 H2O (ℓ) ® 4 H3PO4 (aq)
AlCl3(s) + 6 H2O (ℓ) ® [Al(H2O)6]3+(aq) + 3 Cl-(aq)
SiCl4(l) + 2 H2O(ℓ) ® SiO2(s) + 4 HCl(aq)
Ca3N2 (s) + 6 H2O (ℓ) ® 3 Ca(OH)2 (aq) + 2 NH3(aq)
CaC2 + 2 H2O(ℓ) ® Ca(OH)2 (aq) + C2H2
Al4C3 + 12 H2O(ℓ) ® 4 Al(OH)3 + 3 CH4
4. Formation of hydrates : Water reacts with certain metal salts to form compounds known as hydrates. Depending upon the mode of linkage of water molecules, the following types of hydrates are possible.
· In some hydrates , water molecules act as ligands and get attached to metal ion by co-ordinate bonds thus forming the complex ion e.g.
[Fe(H2O)6]3+ (Cl- )3 , [Ni(H2O)6]2+(NO3- )2
[Cr(H2O)6]3+(Cl- )3
· In some hydrates , water molecule get attached to certain oxygen containing anions through hydrogen bonds. E.g. CuSO4.5H2O. In this hydrate, four water molecules are co-ordinated to central Cu2+ ion , while the fifth water molecule is attached to the sulphate group by hydrogen bonds.
· In some hydrates, water molecules occupy the interstitial sites (voids) in the crystal lattice. BaCl2. 2 H2O is an example of this type of hydrate.
Problem
04. How many hydrogen bonded water molecules are associated in CuSO4 5 H2O ?
HARD AND SOFT WATER
The water obtained from natural source is largely used for washing purposes. The water obtained from a particular source may or may not produce a rich lather with soap generally used for washing purposes. Hence it is important to study the behaviour of water (obtained from a natural source) towards soap solution. On this basis, water can be classified into following two categories.
(i) Soft water : Water which produces lather with soap solution readily is called soft water. Distilled water, rain water etc are some examples of this type of water.
(ii) Hard water : Water which does not produce lather with soap solution readily is called hard water. Sea water, river water, water from a tube well etc. are common examples of hard water.
CAUSE OF HARDNESS
Natural water contains many salts dissolved in it. These salts get dissolved in it , when it passes over rocks or through the various underlying layers of the earth. The most common salts that dissolve in water are bicarbonates, chlorides and sulphates of calcium and magnesium. The presence of these salts makes water hard. Thus hard water may be defined as the water which contains bicrbonates, chlorides and sulphates of calcium and magnesium. Water containing any of these salts does not produce lather with soap readily. This can be explained as follows :
Soaps are sodium or potassium salts of higher fatty acids like palmitic acid (C15H31COOH), stearic acid (C17H35COOH) etc. Thus , the compounds like sodium palmitate (C15H31COONa), sodium stearate (C17H35COONa) etc act as soaps. When hard water containing Ca2+ or Mg2+ ions is treated with a soap solution, a precipitate of the calcium or magnesium salt of the corresponding fatty is obtained as shown below:
Thus , with the use of hard water , soap gets precipitated in the form of calcium and magnesium salts which are insoluble in water. Hence , no lather is produced till all the calcium or magnesium ions get precipitated. This results in more consumption and hence in the wastage of soap. Therefore , hard water is not suitable for washing purposes.
Apart from being unsuitable for laundry washing and dyeing , hard water is harmful for steam boilers. Over a period of time the inner surface of the boiler gets crusted with the so called ‘’boiler scale’’ which is largely calcium sulphate, calcium carbonate and magnesium oxycloride. The deposition reduces the efficiency of the boiler and also damages it. It is therefore necessary to render the water soft before it can be used.
TYPES OF HARDNESS
The hardness of water is of following two types :
1. Temporary hardness : This is due to the presence of bicarbonates of calcium and magnesium dissolved in water. The hardness due to bicarbonates of calcium and magnesium is termed as temporary because it can very easily be removed by simple boiling the hard water.
2. Permanent hardness : This type of hardness is due to the presence of chlorides and sulphates of calcium and magnesium dissolved in water. This cannot be removed easily i.e., just by boiling the hard water and hence called permanent hardness.
SOFTENING OF WATER
The process of removal of Ca2+ and Mg2+ ions from hard water is termed as softening of hard water. There are several methods used to remove both types of hardness. The important methods used for softening of hard water are as follows.
Removal of Temporary hardness
Temporary hardness can be removed by the following methods.
1.By boiling : On boiling the temporary hard water for about fifteen minutes, the bicarbonates of calcium and magnesium present in water decomposes into their insoluble carbonates and carbon dioxide.
The insoluble carbonates thus formed settle at the bottom and can be removed either by filtration or by decantation. The water thus obtained is soft.
2. By Clark’s method : In this method , a calculated amount of milk of lime Ca(OH)2 is added to temporary hard water. This treatment converts the soluble bicarbonates into insoluble carbonates which get precipitated and can be removed by filtration.
REMOVAL OF PERMANENT HARDNESS
Permanent hardness can be removed by any one of the following methods.
1. By adding washing soda : When permanent hard water is treated with calculated amount of washing soda (Na2CO3), the soluble chlorides and sulphates of calcium and magnesium are converted into their insoluble carbonates which get precipitated and may be removed by filtration.
2. Calgon process : A convenient method to soften hard water is to render Ca2+ and Mg2+ ions ineffective instead of removing them from water. This is done by the addition of sodium polymetaphosphate (NaPO3)n , where the value of n is some times large as 1000. However, sodium hexametaphosphate (NaPO3)6 or Na2[Na4(PO3)6] is generally used for this purpose. Calgon is the trade name given to this complex salt.
When hard water containing calcium and magnesium ions is treated with calgon, complex soluble salts Na2[Ca2 (PO3)6] are formed.
The complex salts thus formed remain dissolved in water. They do not cause hardness because the ions which cause hardness (i.e., Ca2+ and Mg2+ ions) are not free. They have been trapped by calgon.
3. Ion exchange method : This is a modern technique widely used to soften hard water. In this technique Ca2+ and Mg2+ ions which cause hardness are exchanged by those cations which do not cause hardness. The exchange of ions is done with the help of certain substances called ion exchangers. Following two types of ion exchangers are generally used for this purpose.
(a) Inorganic cation exchangers(Permutit) : Certain complex inorganic salts possesses the property of exchanging Ca2+ and Mg2+ ions present in hard water for Na+ ions. An important salt of this type is known as sodium aluminium orthosilicate or sodium zeolite (Na2Al2Si2O8. x H2O) . commonly known as permutit. This salt belongs to a family of complex salts known as zeolites. These occur naturally and may also be synthesised . Permutit is an artificially synthesised zeolite.
The permutit is losely packed over a layer of coarse sand in a tank as shown in Fig.
Permutit process for softening of hard water
Hard water is allowed to percolate through it. Ca2+ and Mg2+ ions present in hard water are exchanged with Na+ ions.
As the process continues, permutit gets exhausted due to its conversion into calcium and magnesium zeolites. It takes about 12 hours for permutit to get completely exhausted. Therefore , after a continuous use for 12 hours, it must be regenerated. This is done by percolating a 10% solution of sodium chloride through the exhausted permutit. The operation converts calcium and magnesium zeolites back into sodium zeolite as shown below.
After this treatment, the packing is washed with water to remove the chlorides of cacium and magnesium formed in the above process. The regenerated permutit is now ready for the fresh use.
(b) Organic ion exchangers
Recently some complex organic compounds known as resins have been introduced to remove all minerals from hard water. They remove all cations and anions (except H+ and OH- ions) from hard water and make it completely demineralised.
The resins are very complex organic compounds possessing giant molecules. Two types of resins are used for exchanging all the ions present in hard water. First type of resin which exchanges cations for H+ ions is called cation exchanger. It possesses acidic groups (e.g –COOH or –SO3H group) and may be represented as resin –H+. The second type of resin exchanges anions present in hard water or OH- ions. This is called anion exchanger. It possesses basic groups (e.g. OH- or NH2- ions groups) and can be represented as resin- OH- .
The apparatus used for softening hard water by this method consists of two containers , one containing cation exchange resin and other anion exchange resin over a bed of gravel as shown in Fig.
Softening of water by ion-exchange resins
Hard water is first passed through cation exchange resin. The cations present in hard water get exchanged with H+ ions of the resin.
Ca2+ + 2 resin- H+ ® Ca(resin)2 + 2 H+
Mg2+ + 2 resin- H+ ® Mg(resin)2 + 2 H+
In hard cation
water exchanger
The water coming out of the first container thus contains free H+ ions and is acidic in nature. It is now passed through the anion exchange resin . This resin exchanges the anions with OH- ions.
Cl- + resin- OH- ® resin- Cl- + OH-
SO42- + 2 resin- OH- ® (resin)2-SO42- + 2 OH-
In water Anion exchanger
The OH- ions thus formed combine with free H+ ions (formed in the first container ) to form H2O molecules.
H+ + OH- ® H2O
Thus the water coming out from the second container is neither acidic nor basic and is free from all cations and anions that are present in hard water. It is generally known as deionised or demineralised water.
The continuos use of resins exhaust them after some time. Therefore they must be regenerated. The exhausted resin in the first container is regenerated by percolating a moderately concentrated hydrochloric acid or sulphuric acid solution through it.
Ca(resin)2 + 2 HCl ® 2 resin- H+ + CaCl2
exhausted resin in the
first container
Mg(resin)2 + 2 HCl ® 2 resin- H+ + MgCl2
The exhausted resin in the second container is regenerated by percolating a moderately concentrated sodium hydroxide solution through it.
resin- Cl- + NaOH ® resin- OH- + NaCl
exhausted resin in regenerated resin
second container
This method is widely used for obtaining deionised water for laboratory purposes.
HEAVY WATER (Deuterium oxide , D2O)
Heavy water is the oxide of deuterium. It was discovered by Urey. It occurs to an extent of one part in 6,000 parts of ordinary water.
Preparation of Heavy water
Heavy water is usually prepared by the exhaustive electrolysis of water containing alkali. This method was developed by Taylor, Eyring and Forst in 1933. It involves the electrolysis of N/2 NaOH solution in seven stages using steel cathode and nickel anode.
The electrolytic cell used in this method is shown in Fig. It consists of a cylindrical steel container about 45 cm long and 10 cm in diameter which acts as cathode. A perforated cylindrical sheet at nickel acts as the anode. In actual practice , a number of such cells are generally used.
In the first stage, the electrolysis is carried out for 72 hours when volume of the electrolyte reduces to about one-sixth of the original volume. The alkali present in it is partly neutralised by passing carbon dioxide into it and the contents are distilled.
Electrolytic cell used for the preparation of heavy water
Now, it is taken to another cell for second stage of electrolysis. The process is repeated in seven times. At the end of the seventh stage, the residue consists of 99% D2O. From the third stage onwards, the gases evolved during electrolysis consist of considerable amounts of deuterium. These gases are burnt and the mixture of H2O and D2O thus obtained is added to the electrolyte present in the cell.
Properties of Heavy water
Physical properties : Heavy water is colourless , odourless and tasteless mobile liquid. Almost all the physical constants e.g. specific gravity, melting point, boiling point of heavy water are higher than those of ordinary water. (TABLE Page 5)
Chemical properties
Chemical properties of heavy water are almost similar to those of ordinary water. However, D2O reacts slightly less slowly than H2O. The important chemical properties of D2O are described below.
1. Reaction with active metals : It reacts with active metals like Na, K, Ca etc. to liberate heavy hydrogen.
2 D2O + 2 Na ® 2 NaOD + D2
2 D2O + Ca ® Ca(OD) 2 + D2
2. Electrolysis : When heavy water containing sodium carbonate is subjected to electrolysis, heavy hydrogen is evolved at the cathode.
3. Reaction with metal oxides : It reacts with metal oxides to form corresponding deutroxides. For example,
Na2O + D2O ® 2 NaOD sodium deuteroxide
CaO + D2O ® Ca(OD) 2 calcium deuteroxide
4. Reaction with non-metal oxides : It reacts with the oxides of non-metals to form corresponding deutero acids.
SO3 + D2O ® D2SO4 deuterosulphuric acid
P4O10 + 6 D2O ® 4 D3PO4 deuterophosphoric acid
N2O5 + D2O ® 2 DNO3 deuteronitric acid
5. Reaction with carbides : It reacts with metal carbides to form corresponding heavy hydrocarbons. For example.
CaC2 + 2 D2O ® Ca(OD) 2 + C2D2 deuteroacetylene
Al4C3 + 12 D2O ® 4 Al(OD)3 + 3 CD4 deuteromethane
6. Reaction with nitrides : It reacts with metal nitrides to form deuteroammonia. For example.
Mg3N2 + 6 D2O ® 3 Mg(OD) 2 + 2 ND3 deutroammonia
AlN + 3 D2O ® Al(OD)3 + ND3
7. Formation of deuterohydrates : Heavy water may be associated with salts as water of crystallisation in the same way as the ordinary water does, e.g., Na2SO4 10 D2O ; MgSO4 7 D2O ; CuSO4 5 D2O ; CoCl2 6 D2O etc.,
8. Deuterolysis : Like ordinary water, heavy water causes hydrolysis of certain inorganic salts. The phenomenon is known as deutrolysis. For example,
AlCl3 + 3 D2O ® Al(OD)3 + 3 DCl
Aluminium deterium
deutroxide chloride
9. Exchange reactions : Heavy water reacts with several compounds containing hydrogen to exchange their hydrogen partially or completely by deuterium. Such reactions are termed as exchange reactions. For example,
Biological and physiological effects of heavy water
Heavy water retards the growth of plants. Lewis has shown that tobacco seeds do not germinate in pure heavy water. He further showed that in 50% heavy water the growth of tobacco plant was only half as fast as in pure water. Taylor
Uses of Heavy water
Some important uses of heavy water are as follows :
1. Heavy water is used as a moderator in nuclear reactions involving the fission of uranium, where it slows down the speed of neutrons ejected in the fission process.
2. It is used as a tracer in the study of reactions occurring in living organisms.
3. It is also used for preparation of deuterium.
HYDROGEN ECONOMY (USE OF LIQUID HYDROGEN AS A FUEL)
Man has been prompted to search for alternative and feasible sources of energy because :
· The fossil fuels are limited
· The electricity cannot be stored and
· The use of nuclear power has limitations
A prospective alternative , in this regard , is what is popularly known as ‘’Hydrogen Economy ’’. The basic principle of hydrogen economy is transportation and storage of energy in the form of liquid and gaseous hydrogen.
The main aim and advantage of ‘’hydrogen economy’’ is to transmit energy , not as electrical power , but in the form of hydrogen. This is expected to solve problems associated with the storage and transmission of electricity.
The technology involves the production of bulk quantities of hydrogen electrically and its storage in liquid form in vaccum insulated cryogenic tanks. It is feasible to transport liquid hydrogen by road or rail tankers. It can also be stored in underground tanks and transported by pipe lines. Smaller storage units use metal alloy like LaNi5, Mg-MgH2, Ti-TiH2 etc.
Use of hydrogen as an automobile fuel has many advantages
· Release of greater energy per unit weight of fuel.
· Absence of polluting emissions like CO, CO2, NOX, SO2, hydrocarbons, aldehydes and lead compounds.
· Combustion product is water with just some traces of nitrogen oxides.
· Internal combustion engines can be easily modified for use of hydrogen as fuel.
· Fuel cells for generation of electric power with conversion efficiency of 70-85% have been successfully operated commercially.
HYDROGEN PEROXIDE, H2O2
Hydrogen peroxide was discovered by Thenard , a French scientist in 1818.
Preparation of Hydrogen peroxide
The different methods for the preparation of hydrogen peroxide are discussed below.
1. Laboratory preparation : In the laboratory , hydrogen peroxide can be prepared by any of the following methods.
(a) From sodium peroxide (Merk’s process)
A dilute solution of sulphuric acid (20%) is taken in a container surrounded by ice. Sodium peroxide is now added slowly in small amounts with constant stirring. The solution upon further cooling gives crystals of Na2SO4.10H2O which can be removed by filtration. The solution is 30% aqueous solution of hydrogen peroxide.
Na2O2 + H2SO4 ® Na2SO4 + H2O2
(b) From Barium peroxide : A paste of hydrated barium peroxide (BaO2.8H2O) is prepared in ice cold water and treated with about 20% ice cold solution of sulphuric acid. White precipitate of BaSO4 is removed by filtration and the solution left is about 5% H2O2. It may be noted that anhydrous barium peroxide cannot be used in this method because barium sulphate slowly forms a protective coating around it and prevents it from taking part further in chemical reaction.
BaO2.8H2O + H2SO4 ® BaSO4 + H2O2+ + 8 H2O
white ppt.
The precipitate of BaSO4 formed above is slightly soluble in water. The Ba2+ ions present in solution slowly cause the decomposition of hydrogen peroxide. Therefore the solution cannot be stored for a long time. In order to check this, phosphoric acid is used in place of sulphuric acid. Barium phosphate formed gets completely precipitated and in absence of Ba2+ ions, there is no danger of any decomposition of hydrogen peroxide.
3 BaO2 8 H2O + 2 H3PO4 ® Ba3(PO4)2 +3 H2O2 + 24 H2O
ppt
Hydrogen peroxide can also be obtained by passing a current of CO2 through a cold paste of barium peroxide in water. The white precipitate of barium carbonate is filtered off and the solution is nearly 15 to 20% aqueous solution of hydrogen peroxide.
BaO2 + CO2 + H2O ® BaCO3 +3 H2O2
ppt
2. Industrial preparation : Hydrogen peroxide can be prepared on a commercial scale by the following methods :
(a) By electrolysis of sulphuric acid solution : A 50% solution of sulphuric acid is electrolysed in a cell . As a result , perdisulphuric acid is formed at the anode and hydrogen is evolved at the cathode.
Persulphuric acid is taken from the cell and is then hydrolysed with water to give hydrogen peroxide as follows:
In order to recover hydrogen peroxide, the solution is distilled under reduced pressure . Hydrogen peroxide distils , while sulphuric acid with high boiling point remains undistilled.
The yield of hydrogen peroxide can be improved if a mixture of ammonium sulphate and sulphuric acid taken in equivalent amount is electrolysed.
Ammonium persulphate formed at the anode is withdrawn and is distilled with water to give hydrogen peroxide.
(b) From 2-Ethyl anthraquinone
2-Ethyl anthraquinone is dissolved in benzene and hydrogen gas is passed through the solution in the presence of palladium catalyst when it is reduced to 2-Ethyl anthraquinol.
The product formed is dissolved in a mixture of benzene and cyclohexanol and air is passed. It is oxidised back to 2-ethyl anthraquinone and hydrogen peroxide.
It is an example of alternate reduction and oxidation and is called auto-oxidation.
In this process only oxygen and hydrogen combine in the presence of 2-Ethylanthraquinone to form hydrogen peroxide. This method is quite cheap and is commonly used for the industrial preparation of hydrogen peroxide.
(c) By oxidation of isopropyl alcohol
Isopropyl alcohol is mixed with a small amount of hydrogen peroxide which acts as initiator. Oxygen is passed through the solution at about 340 K and little pressure. As a result , isopropyl alcohol reacts to form acetone and hydrogen peroxide.
Upon distillation acetone and unreacted alcohol distil leaving behind an aqueous solution of hydrogen peroxide.
Concentration of hydrogen peroxide solution
Hydrogen peroxide prepared will be in the form of dilute solution. It cannot be concentrated by boiling because it decomposes below its boiling point to form water and oxygen.
Thus a dilute solution of hydrogen peroxide has to be concentrated very carefully in the following steps.
(i) Evaporation in a water bath : The dilute solution is taken in evaporating dish and is heated carefully in a water bath. Water slowly evaporates and about 50% solution of hydrogen peroxide is obtained. Prolonged heating is not possible because H2O2 will decompose.
(ii) Evaporation in a vaccum desicator : The further concentration is done in a vaccum desicator . The evaporating dish is placed over concentrated sulphuric acid in a vaccum desicator. Upon working the pump , the pressure inside the desicator gets reduced. As a result , water evaporates and vapours are absorbed by the acid (dehydrating agent) . Hydrogen peroxide left in the dish is about 90% concentrated.
(iii) Distillation under reduced pressure : The solution obtained above is distilled under a reduced pressure of 10 to 15 mm. As a result, water distils at a temperature between 303-313 K and about 99% concentrated hydrogen peroxide is left behind.
(iv) Removal of traces of water : In order to remove last traces of water , 99% hydrogen peroxide is surrounded by a freezing mixture of solid carbon dioxide (dry ice) and ether. It gets solidified in the form of crystals. The crystals are removed and dried carefully to remove any water sticking to them. They are then remelted to get completely pure hydrogen peroxide.
Storage of hydrogen peroxide
Hydrogen peroxide readily decomposes. In order to store it , the following precautions must be taken.
(i) It should be kept in polyethene or teflon bottles because glass surface being rough can cause its decomposition. In case, glass bottles are used they must be coated with wax. In order to avoid the decomposition by light, the bottles must be coloured.
(ii) Small amounts of stabilisers like phosphoric acid , glycerol or acetanilide must be added to hydrogen peroxide. They also retard its decomposition and are known as inhibitors or negative catalysts.
Expressing strength of hydrogen peroxide
The strength of a dilute solution of hydrogen peroxide is expressed as percent strength and also as volume strength.
(a) Percent strength : It represents the amount of H2O2 by mass present in 100mL of the solution. For example, a 20% solution means that 20 g of hydrogen peroxide are present in 100 mL of the solution.
(b) Volume strength : The strength of hydrogen peroxide is normally expressed in terms of volume. The volume strength refers to volume of oxygen (measured at STP ) obtained when one volume of the given sample of hydrogen peroxide will give when it is subjected to heating. The bottles are marked as 10 volume, 20 volume or 30 volume etc. depending upon the strength of hydrogen peroxide. A 20 volume solution of hydrogen peroxide means that 1 mL of hydrogen peroxide solution at STP will give 20 mL of oxygen.
A 30% solution of hydrogen peroxide is marketed as ‘100 volume ‘ hydrogen peroxide. Commercially , it is marketed as 10 V, which means it contains 3% H2O2.
Problem
05. Convert 10 volume hydrogen peroxide into percent strength and also as strength in g/L
06. Calculate the normality of 30 volume solution of hydrogen peroxide.
PROPERTIES OF HYDROGEN PEROXIDE
Physical properties
(i) In pure form hydrogen peroxide is a thick syrupy liquid and has a pale blue colour.
(ii) It has a bitter taste.
(iii) Hydrogen peroxide has a density of 1.44 g cm-3 more than that of water(1.0 g cm-3) . The molecules of H2O2 are more associated because of hydrogen bonding than the water molecules.
(iv) Its melting point is 272.4 K. But its boiling point cannot be determined correctly since it decomposes before that temperature. However, the boiling point has been found approximately 423.2 K by extrapolation method.
(v) It is miscible with water , alcohol and ether in all proportions because of hydrogen bonding.
Chemical properties
1. Decomposition : Pure hydrogen peroxide is unstable in nature and decomposes to form water and oxygen upon standing or upon heating accompanied by evolution of heat.
2 H2O2 ® 2 H2O + O2 : DH = - 196.0 kJ
The decomposition can be accelerated in the presence of certain metals(Cu, Ag, Pt etc), metal oxides (MnO2) or even metal ions (Ba2+, Fe2+ etc.). Rough surface of glass can also cause its decomposition. That is why , it is kept in polythene bottles.
2. Acidic character : Pure hydrogen peroxide is weakly acidic in nature and turns blue litmus red. It is slightly stronger acid (Ka = 1.55 x 10-12 at 298 K) than water (Ka = 1.0 x 10-14 at 298 K). As it has two replaceable hydrogen atoms, therefore, it forms two series of salts with base like NaOH.
It may be noted that a dilute solution of H2O2 is almost neutral.
3. Oxidising nature : Since hydrogen peroxide readily decomposes to evolve oxygen, it acts as an oxidising agent. The oxidation take place in neutral, acidic and also in alkaline medium.
(a) In neutral medium : Some of the oxidising reactions are as follows :
(i) It oxidises sulphites to sulphates.
H2O2 ® H2O + O
Na2SO3 + O ® Na2SO4
__________________________
H2O2 + Na2SO3 ® Na2SO4 + H2O
(ii) It oxidises hydrogen sulphide to sulphur.
H2O2 ® H2O + O
H2S + O ® H2O + S
_____________________
H2O2 + H2S ® 2 H2O + S
(iii) It oxidises lead sulphide to lead sulphate.
H2O2 ® H2O + O
PbS + 4[O] ® PbSO4
_____________________
PbS + 4 H2O2 ® PbSO4+ 4 H2O
(black) (white)
Traces of hydrogen sulphide are always present in the atmosphere. They slowly blacken the lead painting containing PbO due to formation of lead sulphide. The colour of the painting can be restored by dipping them in aqueous solution of hydrogen peroxide which converts lead sulphide to lead sulphate (white in colour).
(b) In acidic medium
In presence of acid , H2O2 can accept electrons and act as oxidising agent.
H2O2 + 2 H+ + 2 e- ® 2 H2O
(i) It evolves iodine from acidified KI solution.
H2O2 ® H2O + O
2 KI + H2O + [O] ® K2SO4 + H2O + I2
___________________________________
2 KI + H2SO4 + H2O2 ® K2SO4+ 2 H2O + I2
(ii) It oxidises acidified ferrous sulphate to ferric sulphate.
H2O2 ® H2O + O
2 FeSO4 + H2SO4 + O ® Fe2(SO4)3 + H2O
___________________________________
2 FeSO4 + H2SO4 + H2O2 ® Fe2(SO4)3 + 2 H2O
(iii) It oxidises acidified potassium ferrocyanide to
potassium ferricyanide.
H2O2 ® H2O + O
2 K4Fe(CN)6 + H2SO4 + O ® 2 K3Fe (CN)6 + K2SO4 +H2O
___________________________________________
2 K4Fe(CN)6 + H2SO4 + H2O2 ® 2 K3Fe (CN)6 + K2SO4 + H2O
(iv) When dissolved in ether , hydrogen peroxide oxidises an ice cold solution of acidified potassium dichromate to form chromium pentoxide.
[ H2O2 ® H2O + O ] x 4
2 K2Cr2O7 + H2SO4 +4 [O] ® K2S O4 + 2 CrO5 +H2O
________________________________________
K2Cr2O7+ H2SO4 + 4 H2O2 ® K2S O4 + 2 CrO5 + 5 H2O
chromium pentoxide
The ether layer containing CrO5 become blue.
(c) In alkaline medium
Hydrogen peroxide can also act as oxidising agent in the alkaline medium by accepting electrons.
H2O2 + 2 e- ® 2 OH-
(i) It oxidises manganese sulphate to manganese dioxide in presence of alkali.
MnSO4 + H2O2 + 2 NaOH ® MnO2 + Na2SO4 + 2 H2O
Or Mn2+ + H2O2 + 2 OH- ® MnO2 + 2 H2O (ionic equation)
(ii) It oxidises chromium salts to chromates in alkaline medium.
Cr2(SO4)3 +3 H2O2 +10 NaOH ® 2 Na2CrO4 +3 Na2SO4+ 8 H2O
Sodium chromate
2 Cr3+ + 3 H2O2 + 10 OH- ® 2 CrO42- + 8 H2O (ionic equation)
4. Reducing nature
Hydrogen peroxide can also take up oxygen from strong oxidising agents and can act as reducing agent.
H2O2 + O ® H2O + O2
(a) In neutral medium
(i) Halogens are reduced to halogen acids.
Cl2+ H2O2 ® 2 HCl + O2
It can remove the unreacted chlorine used in the process of bleaching. Therefore it is called antichlor.
(ii) Oxides of metal are reduced to metals.
Ag2O + H2O2 ® 2 Ag + H2O + O2
PbO2 + H2O2 ® PbO + H2O + O2
(b) In alkaline medium
The reaction takes place in the presence of a strong alkali such as NaOH or KOH. For example, potassium ferricyanide is reduced to potassium ferrocyanide.
2 K3[Fe(CN)6] + 2 KOH ® 2 K4[Fe(CN)6] + 2 H2O + O2
H2O2 + O ® H2O + O2
2 K3[Fe(CN)6]+2 KOH+ H2O2 ®2 K4[Fe(CN)6 +2 H2O + O2
or 2 [Fe(CN)6]2-+ 2 KOH + H2O2 ® 2 Fe[CN)6]4-+2 H2O + O2
ferricyanide ferrocyanide
5. Bleaching action
Hydrogen peroxide acts as a bleaching agent due to oxidation. It oxidises the colouring matter to the colouless product.
H2O2 ® H2O + O
Colouring matter + O ® colourless mass (bleached)
It is used to bleach fine materials such as ivory , silk, wool, hair etc.
STRUCTURE OF HYDROGEN PEROXIDE
Hydrogen peroxide is a dihydroxy compound ( H- O - O- H ) and the O- O linkage is known as peroxide linkage. It is a non-planar molecule as the two O - H bonds are in different planes. The interplanar (dihedral) angle is 111.5° in the gaseous phase but is reduced to 90.2° in the crystalline state because of hydrogen bonding. The molecular dimensions in solid and gaseous phases are shown in the Fig.
(a) Gas phase
(b) Solid phase
Fig. (a) H2O2 structure (gas phase) . The dihedral angle 111.5° (b) H2O2 (solid phase at 110 K) The dihedral angle is reduced to 90.2°. The O- O distance is of a normal single bond.
Uses of hydrogen peroxide
1. The largest industrial use of H2O2 is as a bleach for textiles, paper pulp, straw, leather, oils, fats, etc.
2. Domestically it is used as a hair bleach and as a mild disinfectant.
3. It is extensively used for the manufacture of chemicals like sodiumperborate and percarbonate which are important constituents of high quality detergents.
4. H2O2 is used for the production of epoxides, propylene oxide and polyurethanes.
5. It is used for the synthesis of hydroquinone, pharmaceuticals (cehalosoporin) and food products like tartaric acid.
6. A fast growing use of H2O2 is in environmental (Green) chemistry, for example in pollution control treatment of domestic and industrial effluents ; oxidation of cyanides and restoration of aerobic conditions to sewage wastes. This use has led to tremendous increase in the industrial production of H2O2 .
Tests for Hydrogen peroxide
1. It decolorises acidified potassium permanganate solution.
2. When an etherial solution of H2O2 is shaken with acidified potassium dichromate solution, the etherial layer acquires a blue colour due to the formation of CrO5.
3. It liberates iodine from acidified KI solution which turns starch paper blue.
4. It gives an orange colour with the solution of titanium dioxide (TiO2) due to the formation of pertitanic acid (H2TiO4)
QUESTIONS
1. Justify the position of hydrogen in the periodic table on the basis of its electronic configuration.
2. Write the names of the three isotopes of hydrogen. What is the mass ratio of these isotopes ?
3. Why does hydrogen occur in diatomic form rather than in monoatomic form under normal conditions.
4. How can the production of dihydrogen from ‘coal gasification’ be increased ?
5. Describe the bulk preparation of dihydrogen by electrolytic method. What is the role of an electrolyte in the process ?
6. Complete the following :
7. Describe the consequences of high bond enthalpy of H – H bond in terms of chemical reactivity of dihydrogen.
8. What do you understand by (i) electron deficient (ii) electron precise and (iii) electron rich compounds of hydrogen ? Provide justification with suitable example.
9. What characteristics do you expect from electron deficient hydride with respect to its structure and chemical reaction ?
10. Do you expect the carbon hydride of the type (CnH2n+2 ) to act as Lewis acid or base ?
11. What do you understand by the term non-stiochiometric hydrides ? Do you expect these types of hydrides to be formed from alkali metals ? Justify your answer.
12. How would you expect the metallic hydrides to be useful for hydrogen storage ? Explain.
13. How does atomic hydrogen or oxy-hydrogen torch function for cutting and welding purposes. Explain.
14. Among NH3, H2O and HF which would you expect to have highest magnitude of hydrogen bonding and why ?
15. Saline hydrides are known to react violently water producing fire. Can CO2 a well known extinguisher , be used in this case ? Explain.
16. Arrange the following :
(i) LiH , NaH and CsH in order of increasing ionic character.
(ii) H – H , D – D and F – F in order of increasing bond dissociation enthalpy.
(iii) NaH, MgH2 and H2O in order of increasing reducing power.
17. Compare the structure of H2O and H2O2.
18. What dou you understand by ‘auto protolysis’ of water ? What is its significance ?
19. Complete the following :
(i) Pb(s) + H2O2 (aq) ®
(ii) MnO4- + (aq) + H2O2 (aq) ®
(iii) CaO(s) + H2O2 (aq) ®
(iv) AlCl3(s) + H2O2 (aq) ®
(v) Ca3N2 + H2O2 (aq) ®
20. Consider the reaction of water with fluorine and suggest in terms of oxidation and reduction which species are oxidized/reduced.
21. Discuss the structure of ice.
22. What causes temporary and permanent hardness of water ?
23. Discuss the principle and method of softening of hard water by synthetic ion-exchange method.
24. Write chemical reactions to show amphoteric nature of water.
25. Hydrogen peroxide acts both as oxidizing and as reducing agent. Explain.
26. Explain why hydrogen is best placed separately in the periodic table of elements ?
27. What is understood by ‘’Water gas shift reaction’’ ?
28. Hydrogen forms compounds with elements having atomic numbers : 9, 11, 12 and 17. What are their chemical behaviour.
29. What are metallic/interstitial hydrides ? How do they differ from molecular hydrides ?
30. Complete the following reactions :
(i) CaO + H2O ® ? (ii) Na2O (s) + H2O ® ?
31. Compare the structures of H2O and H2O2.
32. Explain why hydrogen peroxide is stored in coloured /plastic bottles ?
33. Describe the industrial applications of hydrogen dependent on :
(i) the heat liberated when atoms are made to combine on the surface of a metal.
(ii) Its effect on unsaturated organic systemsin presence of a catalyst.
(iii) Its ability to combine with nitrogen under specific conditions.
34. How is dihydrogen prepared :
(i) from water by using a reducing agent ?
(ii) in laboratory in pure form ?
(iii) from hydrocarbons ?
35. Complete the following equations :
(i) Fe(s) + H2O(g) ® ?
(ii) PbS(s) + H2O2(aq) ® ?
(iii) MnO4-+ H2O2(aq) ® ?
36. Discuss the importance of heavy water in nuclear reactors ?
37. How is heavy water prepared from normal water ?
38. Explain why water has high boiling point and melting point as compared to H2S ?
39. Describe the structure of common form of ice.
40. Distinguish between (i) hard and soft water (ii) temporary and permanent hardness.
41. Discuss the principle and method of sofening of hard water by organic ion exchange resins.
42. Explain amphoteric nature of water.
43. Show by proper chemical reactions how hydrogen peroxide can function both as oxidising and reducing agent ?
44. What is understood by hydrogen economy ?
45. Explain the correct context in which the following terms are used :
(i) diprotium (ii) dihydrogen (iii) proton (iv) hydrogen
46. Is it correct to say that hydrogen can behave as a metal ? State the conditions under which such behaviour can be possible.
47. Name the isotopes of hydrogen. What is the importance of heavier isotopes of hydrogen.
48. How many allotropes of dihydrogen are known ? What is their importance ?
49. Name the class of hydrides to which H2O, B2H6 , NaH and LaH3 belong ? What is understood by ‘Hydride Gap’ ?
50. Hydrogen forms three types of bonds in its compounds. Describe each type of bonding using suitable examples.
51. Elements with atomic number 17 and 20 form compounds with hydrogen. Write the formulae of these two compounds and compare their chemical behaviour in water.
52. Give an example each of an ionic hydride and a covalent hydride.
53. Complete the following reactions :
(i) CuO(s) + H2 (g) ® ?
(ii) CO(g) + H2(g) ® ?
54. Compare the chemical properties of H2O and H2O2.
55. Describe some unusual properties of water.
56. What is the difference between hydrolysis and hydration ?
57. What is understood by hydrogenation ?
58. What are the advantages in using hydrogen as a fuel ?
59. Ionic hydrides are frequently used to remove traces of water from organic compounds. What is the underlying basis of this process ?
60. Although D2O resembles H2O chemically yet it is a toxic substance. Explain.
61. Why do lakes freeze from the top towards bottom ?
62. Why is ice less dense than water and what kind of attractive forces must be overcome to melt ice ?
63. Name radioactive isotope of hydrogen.
64. How will you prepare a sample of ND3 ?
65. Why is dihydrogen gas not preferred in balloons ?