Hydrogen : Class 11
Hydrogen: Class 11
hydrogen
Position of Hydrogen in the Periodic Table:
Hydrogen is similar to both G1 (Alkali metals) and G17 (Halogens) elements in the periodic table. Hence it is not right to keep H either in G1 or G17 of the periodic table.
Similarity with Alkali metals and Halogens:
Hydrogen is similar to alkali metals, which lose one electron to form unipositive ions.
Also, hydrogen is similar to halogens which gain one electron to form uninegative ions.
Dissimilarity with Alkali metals and Halogens
Unlike alkali metals, H has a very high ionization enthalpy and does not possess metallic characteristics under normal conditions.
In terms of reactivity, it is very low as compared to halogens.
Isotopes of hydrogen:
Ø 1H1 : Protium , Most abundant in nature
Ø 2H1: Deuterium (D), Component of heavy water.
Ø 3H1 : Tritium (T) , Radioactive in nature
Dihydrogen:
a. Laboratory preparation:
Reaction of metals with acids: Zn + H+→ Zn2++H2
b. Commercial Preparation:
1. Electrolysis of acidified water
2. Electrolysis of warm aqueous Ba(OH)2 between nickel electrodes. (for the preparation of high purity (>99.95%) dihydrogen)
3. By-product in the manufacture of NaOH and Cl2 by electrolysis of brine solution.
4. Bosch’s Process
The maximum quantity of commercial dihydrogen is prepared by this method. It involves the following steps:
(a) Preparation of water gas.
It is prepared by the action of steam with red hot coke
Water gas can also be obtained by action of steam on hydrocarbons in the presence of catalyst. This process is also called steal reforming of hydrocarbons
(b) Separation of Hydrogen.
The water gas formed by any of the above methods, is mixed with steam and is passed over heated Fe2O3 and Cr2O3 at 770K when CO is oxidized to CO2. Gaseous mixture of CO2 and H2 is then bubbled into cold water under pressure when CO2 dissolves leaving behind dihydorgen gas which escapes out.
This reaction is called water gas-shift reaction.
c. Properties:
Ø Reaction with halogen:
H2 + X2→ 2HX [X= F, Cl, Br, I]
Ø Reaction with oxygen:
H2 (g) +O2 (g) + Δ→ 2H2O(l);
ΔH0 = -285.9 kJmol-1
Ø Reaction with nitrogen:
H2(g)+N2(g)+Δ→2NH3(l) ; ΔH0=-92kJ mol-1
Ø Reaction with alkali metals:
3H2(g) +2M(g) + Δ → 2MH(s)
Ø Reaction with metal oxide
Reduces such as iron and metals less active than iron.
Ø Reaction with carbon monoxide
Dihydrogen reacts with carbon monoxide at 700 K in the presence of a catalyst Zn/Cr2O3 to produce methanol.
Ø Reaction with unsaturated hydrocarbons:
Unsaturated hydrocarbons like ethene and ethyne react with dihydrogen to form saturated hydrocarbons.
Uses of Hydrogen:
ü Used for synthesis of ammonia and vanaspati fat and many other products.
ü Used as rocket fuel.
ü Used in hydrogen fuel cells.
Hydrides
Dihydrogen combines with a number of elements to form binary compounds called hydrides. Their general formula being MHx where M represents the element and x the number of hydrogen atoms
1. Ionic or salt-like or saline hydrides
2. Metallic or Interstitial hydrides
3. Molecules or Covalent hydrides
1. Saline Hydrides or Ionic hydrides:
These are binary compounds of hydrogen and elements which are more electropositive than hydrogen such as alkali metals, alkaline earth metals (except Be), etc. Some common examples of this category are: LiH, NaH, CaH2, CrH2 etc.
The general characteristic of these hydrides are as follows:
(i) They are crystalline solids having white or greyish colour.
(ii) They have high melting and boiling points.
(iii) They have high density and high heat of formation.
(iv) They conduct electricity in molten state liberating dihydrogen gas at anode which confirm the presence of hydride (H-) in them.
CaH2 (melt) → Ca2+ + 2H-
At anode: 2H- → H2 + 2e-
At cathode: Ca2+ + 2e- → Ca
2. Covalent Hydrides or Molecular Hydrides:
These are the compounds of hydrogen formed with most of the p-block elements. The sub-types are :
ü Electron deficient hydrides:- The hydrides which do not have sufficient number of electrons to form normal covalent bonds is called electron deficient hydride. For example, hydride of group 13 (BH3, AlH3, etc.).They are known as Lewis acids i.e., electron acceptors. To make up their deficiency they generally exist in polymeric forms such as B2H6, Al2H6, etc.
ü Electron precise hydrides:-The hydrides which have sufficient number of electrons required for forming covalent bonds is called electron precise hydride. For example, hydrides of group 14 (CH4, SiH4, GeH4, SnH4, PbH4 etc.) they have tetrahedral geometry.
ü Electron rich hydrides:-The hydrides which have excess electrons as required to form normal covalent bonds is called electron rich hydride. For example, hydrides of group 15 to 17 (NH3, PH3, H2O, H2S, H2Se, H2Te, HF etc.)
3. Interstitial Hydride or Metallic Hydrides:
These are binary compounds of hydrogen and transition elements.
These hydrides are generally formed by the
(a) transition metals of group 3, 4, 5 of d- block;
(b) Cr metal of group 6 and
(c) f – block elements.
It may be noted that elements of group 7, 8, 9 of d – block do not form hydrides at all. This inability of metal, of group 7, 8, 9 of periodic table to form hydrides is referred to as hydride gap of d – block.
In these compounds H atoms are supposed to occupy interstitial position in the metal lattices. The composition of these hydrides may not correspond to simple whole number ratio and therefore, they are also called non-stoichimotric hydrides. Some examples of interestial hydrides of elements of group 3 to 5 are ScH2 , YH2 ,YH3 ,LiH3, CrH, TiH2,ZrH2, HfH2 etc.
Some examples of non-stoichimetric hydrides are PdH0.6, TiH1.8, etc.
Water: -
It is the major part of all living organisms.water is also known as the river of life. Human body has about 65%and some plants have as much as 95%water.
Structure Of Water:- In a gas phase water is bent molecule with a bond angle of 104.5 and O-H bond length of 95.7pm. It is highly polar molecule.
Structure of ice:-Ice has a highly ordered 3D hydrogen bonded structure. Each oxygen atom is surrounded tetrahedrally by four other four other oxygen atoms at a distance of 276 pm.
· Properties of Water
Physical properties
Water has some unique features which arise due to intermolecular H-bonding.
(i) Water because of its high dielectric constant (78.39) has the ability to dissolve most of the inorganic (ionic) compounds and its, therefore, regarded as a universal solvent. Whereas solubility of ionic compounds takes place due to ion-dipole interactions (i.e. solvation of ions), the solubility of covalent compounds such as alcohols, amines, urea, glucose, sugar etc, takes place due to the tendency of these molecules to form hydrogen bonds with water.
(i) The freezing point, boiling point, heat of fusion and heat of vapourization water are abnormally higher than those of the hydrides of the other elements of the same group (16) such as H2S, H2Se, H2Te etc. This is due to the presence of intermolecular hydrogen bonding in H2O molecules.
Chemical Properties of Water
1. Action towards litmus
Pure water is neutral to litmus.
2. Decomposition
Water is quite stable and does not dissociate at high temperature. The dissociation into its elements is only 0.02 % even at 1500K.
3. Acid-base reactions
Water is amphoteric substance because it can act as acid as well as base as shown -
However, the pH of water at 250C is 7 and it is neutral towards litmus.
4. Hydrolytic reactions
Water can hydrolyse many non-metallic oxides, halides and also some metallic phosphides, carbides and nitrates.
Hard Water:
Water containing carbonate, chloride and sulphate salts of calcium and magnesium is said to be hard.
Ø Temporary hardness is due to the presence of bicarbonate salts and can be removed by boiling or by adding lime water.
· Ca(HCO3)2 + Ca(OH)2 → 2CaCO3↓+2H2O (Clark’s method)
· Ca(HCO3)2 + Δ → 2CaCO3↓+2H2O +CO2↑
·
Ø Permanent hardness is due to presence of sulphate and chloride salts and can be removed by treatment with washing soda o.
· MCl2 + Na2CO3 → MCO3 ↓ +2NaCl (M= Mg, Ca)
· MSO4 + Na2CO3 → MCO3↓ +Na2SO4 (M= Mg, Ca)
Ø Hard water forms scum/precipitate with soap:
Heavy water:
· Molecular formula: D2O
· 10.68% denser than ordinary water
· Freezing point 3.8 0C
· Unfit for drinking and causes sterility.
· It is used as a moderator in nuclear reactors to speed down the nuclear reactions.
Hydrogen peroxide :
a. Preparation :
Ø Lab Method:
· Na2O2(s)+ H2SO4(aq) → H2O2(aq) + Na2SO4(s)
· BaO2.8H2O+H2SO4(aq)→H2O2(aq) + BaSO4(s)
· Anhydrous barium oxide is not used because the precipitated BaSO4 forms a protective layer on the unreacted barium peroxide and thus prevents its further participation in the reaction. However it can be overcome by using phosphoric acid.
Ø By Electrolysis:
Ø By the auto-oxidation of 2-ethyl anthraquinol.
The net reaction is a catalytic union of H2 and O2 to yield hydrogen peroxide.
b. Properties
i) Unstable liquid, decomposes to give water and dioxygen and the reaction is slow in the absence of catalyst. It is catalysed by certain metal ions, metal powders and metal oxides.
2H2O2 (l) → 2H2O (l) + O2 (g)
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 V of oxygen at STP. Commercially, it is marketed as 10 V, which means it contains 3% H2O2.
ii) It is a very powerful oxidising agent and poor reducing agent.
As oxidising agent
In acidic medium: H2O2 + 2H+ + 2e- → 2H2O
In basic medium :H2O2 + OH- + 2e- → 3OH--
As reducing agent
In acidic medium: H2O2 → 2H+ + O2 + 2e-
In basic medium : H2O2 + 2OH- → 2H2O + O2 + 2e-
2Fe2+ + H2O2 + 2H+ → 2Fe3+ + 2H2O
2MnO4- + 5H2O2 + 6H+ → 2Mn2+ + 8H2O + 5O2
Mn2+ + H2O2 → Mn+4 + 2OH-
2Fe3+ + H2O2 + 2OH- → 2Fe2+ + 2H2O + O2
The oxidising property of hydrogen peroxide is put to use in the restoration of old paintings, where the original white lead paint has been converted to black PbS by the H2S in the atmosphere. Hydrogen peroxide oxidises the black PbS into white PbSO4.
PbS(s) + 4H2O2 (aq) → PbSO4(s) + 4H2O
black white
c. Tests :
· It liberates iodine from potassium iodide in presence of ferrous sulphate
· Acidified solution of dichromate ion forms a deep blue colour with H2O2 due to the formation of CrO5. ,
Cr2O72- + 4H2O2 + 2H+ → 2CrO5 +5H2O
· With a solution of titanium oxide in conc.H2SO4, it gives orange colour due to the formation of pertitanic acid.
Position of Hydrogen in the Periodic Table:
Hydrogen is similar to both G1 (Alkali metals) and G17 (Halogens) elements in the periodic table. Hence it is not right to keep H either in G1 or G17 of the periodic table.
Similarity with Alkali metals and Halogens:
Hydrogen is similar to alkali metals, which lose one electron to form unipositive ions.
Also, hydrogen is similar to halogens which gain one electron to form uninegative ions.
Dissimilarity with Alkali metals and Halogens
Unlike alkali metals, H has a very high ionization enthalpy and does not possess metallic characteristics under normal conditions.
In terms of reactivity, it is very low as compared to halogens.
Isotopes of hydrogen:
Ø 1H1 : Protium , Most abundant in nature
Ø 2H1: Deuterium (D), Component of heavy water.
Ø 3H1 : Tritium (T) , Radioactive in nature
Dihydrogen:
a. Laboratory preparation:
Reaction of metals with acids: Zn + H+→ Zn2++H2
b. Commercial Preparation:
1. Electrolysis of acidified water
2. Electrolysis of warm aqueous Ba(OH)2 between nickel electrodes. (for the preparation of high purity (>99.95%) dihydrogen)
3. By-product in the manufacture of NaOH and Cl2 by electrolysis of brine solution.
4. Bosch’s Process
The maximum quantity of commercial dihydrogen is prepared by this method. It involves the following steps:
(a) Preparation of water gas.
It is prepared by the action of steam with red hot coke
Water gas can also be obtained by action of steam on hydrocarbons in the presence of catalyst. This process is also called steal reforming of hydrocarbons
(b) Separation of Hydrogen.
The water gas formed by any of the above methods, is mixed with steam and is passed over heated Fe2O3 and Cr2O3 at 770K when CO is oxidized to CO2. Gaseous mixture of CO2 and H2 is then bubbled into cold water under pressure when CO2 dissolves leaving behind dihydorgen gas which escapes out.
This reaction is called water gas-shift reaction.
c. Properties:
Ø Reaction with halogen:
H2 + X2→ 2HX [X= F, Cl, Br, I]
Ø Reaction with oxygen:
H2 (g) +O2 (g) + Δ→ 2H2O(l);
ΔH0 = -285.9 kJmol-1
Ø Reaction with nitrogen:
H2(g)+N2(g)+Δ→2NH3(l) ; ΔH0=-92kJ mol-1
Ø Reaction with alkali metals:
3H2(g) +2M(g) + Δ → 2MH(s)
Ø Reaction with metal oxide
Reduces such as iron and metals less active than iron.
Ø Reaction with carbon monoxide
Dihydrogen reacts with carbon monoxide at 700 K in the presence of a catalyst Zn/Cr2O3 to produce methanol.
Ø Reaction with unsaturated hydrocarbons:
Unsaturated hydrocarbons like ethene and ethyne react with dihydrogen to form saturated hydrocarbons.
Uses of Hydrogen:
ü Used for synthesis of ammonia and vanaspati fat and many other products.
ü Used as rocket fuel.
ü Used in hydrogen fuel cells.
Hydrides
Dihydrogen combines with a number of elements to form binary compounds called hydrides. Their general formula being MHx where M represents the element and x the number of hydrogen atoms
1. Ionic or salt-like or saline hydrides
2. Metallic or Interstitial hydrides
3. Molecules or Covalent hydrides
1. Saline Hydrides or Ionic hydrides:
These are binary compounds of hydrogen and elements which are more electropositive than hydrogen such as alkali metals, alkaline earth metals (except Be), etc. Some common examples of this category are: LiH, NaH, CaH2, CrH2 etc.
The general characteristic of these hydrides are as follows:
(i) They are crystalline solids having white or greyish colour.
(ii) They have high melting and boiling points.
(iii) They have high density and high heat of formation.
(iv) They conduct electricity in molten state liberating dihydrogen gas at anode which confirm the presence of hydride (H-) in them.
CaH2 (melt) → Ca2+ + 2H-
At anode: 2H- → H2 + 2e-
At cathode: Ca2+ + 2e- → Ca
2. Covalent Hydrides or Molecular Hydrides:
These are the compounds of hydrogen formed with most of the p-block elements. The sub-types are :
ü Electron deficient hydrides:- The hydrides which do not have sufficient number of electrons to form normal covalent bonds is called electron deficient hydride. For example, hydride of group 13 (BH3, AlH3, etc.).They are known as Lewis acids i.e., electron acceptors. To make up their deficiency they generally exist in polymeric forms such as B2H6, Al2H6, etc.
ü Electron precise hydrides:-The hydrides which have sufficient number of electrons required for forming covalent bonds is called electron precise hydride. For example, hydrides of group 14 (CH4, SiH4, GeH4, SnH4, PbH4 etc.) they have tetrahedral geometry.
ü Electron rich hydrides:-The hydrides which have excess electrons as required to form normal covalent bonds is called electron rich hydride. For example, hydrides of group 15 to 17 (NH3, PH3, H2O, H2S, H2Se, H2Te, HF etc.)
3. Interstitial Hydride or Metallic Hydrides:
These are binary compounds of hydrogen and transition elements.
These hydrides are generally formed by the
(a) transition metals of group 3, 4, 5 of d- block;
(b) Cr metal of group 6 and
(c) f – block elements.
It may be noted that elements of group 7, 8, 9 of d – block do not form hydrides at all. This inability of metal, of group 7, 8, 9 of periodic table to form hydrides is referred to as hydride gap of d – block.
In these compounds H atoms are supposed to occupy interstitial position in the metal lattices. The composition of these hydrides may not correspond to simple whole number ratio and therefore, they are also called non-stoichimotric hydrides. Some examples of interestial hydrides of elements of group 3 to 5 are ScH2 , YH2 ,YH3 ,LiH3, CrH, TiH2,ZrH2, HfH2 etc.
Some examples of non-stoichimetric hydrides are PdH0.6, TiH1.8, etc.
Water: -
It is the major part of all living organisms.water is also known as the river of life. Human body has about 65%and some plants have as much as 95%water.
Structure Of Water:- In a gas phase water is bent molecule with a bond angle of 104.5 and O-H bond length of 95.7pm. It is highly polar molecule.
Structure of ice:-Ice has a highly ordered 3D hydrogen bonded structure. Each oxygen atom is surrounded tetrahedrally by four other four other oxygen atoms at a distance of 276 pm.
· Properties of Water
Physical properties
Water has some unique features which arise due to intermolecular H-bonding.
(i) Water because of its high dielectric constant (78.39) has the ability to dissolve most of the inorganic (ionic) compounds and its, therefore, regarded as a universal solvent. Whereas solubility of ionic compounds takes place due to ion-dipole interactions (i.e. solvation of ions), the solubility of covalent compounds such as alcohols, amines, urea, glucose, sugar etc, takes place due to the tendency of these molecules to form hydrogen bonds with water.
(i) The freezing point, boiling point, heat of fusion and heat of vapourization water are abnormally higher than those of the hydrides of the other elements of the same group (16) such as H2S, H2Se, H2Te etc. This is due to the presence of intermolecular hydrogen bonding in H2O molecules.
Chemical Properties of Water
1. Action towards litmus
Pure water is neutral to litmus.
2. Decomposition
Water is quite stable and does not dissociate at high temperature. The dissociation into its elements is only 0.02 % even at 1500K.
3. Acid-base reactions
Water is amphoteric substance because it can act as acid as well as base as shown -
However, the pH of water at 250C is 7 and it is neutral towards litmus.
4. Hydrolytic reactions
Water can hydrolyse many non-metallic oxides, halides and also some metallic phosphides, carbides and nitrates.
Hard Water:
Water containing carbonate, chloride and sulphate salts of calcium and magnesium is said to be hard.
Ø Temporary hardness is due to the presence of bicarbonate salts and can be removed by boiling or by adding lime water.
· Ca(HCO3)2 + Ca(OH)2 → 2CaCO3↓+2H2O (Clark’s method)
· Ca(HCO3)2 + Δ → 2CaCO3↓+2H2O +CO2↑
·
Ø Permanent hardness is due to presence of sulphate and chloride salts and can be removed by treatment with washing soda o.
· MCl2 + Na2CO3 → MCO3 ↓ +2NaCl (M= Mg, Ca)
· MSO4 + Na2CO3 → MCO3↓ +Na2SO4 (M= Mg, Ca)
Ø Hard water forms scum/precipitate with soap:
Heavy water:
· Molecular formula: D2O
· 10.68% denser than ordinary water
· Freezing point 3.8 0C
· Unfit for drinking and causes sterility.
· It is used as a moderator in nuclear reactors to speed down the nuclear reactions.
Hydrogen peroxide :
a. Preparation :
Ø Lab Method:
· Na2O2(s)+ H2SO4(aq) → H2O2(aq) + Na2SO4(s)
· BaO2.8H2O+H2SO4(aq)→H2O2(aq) + BaSO4(s)
· Anhydrous barium oxide is not used because the precipitated BaSO4 forms a protective layer on the unreacted barium peroxide and thus prevents its further participation in the reaction. However it can be overcome by using phosphoric acid.
Ø By Electrolysis:
Ø By the auto-oxidation of 2-ethyl anthraquinol.
The net reaction is a catalytic union of H2 and O2 to yield hydrogen peroxide.
b. Properties
i) Unstable liquid, decomposes to give water and dioxygen and the reaction is slow in the absence of catalyst. It is catalysed by certain metal ions, metal powders and metal oxides.
2H2O2 (l) → 2H2O (l) + O2 (g)
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 V of oxygen at STP. Commercially, it is marketed as 10 V, which means it contains 3% H2O2.
ii) It is a very powerful oxidising agent and poor reducing agent.
As oxidising agent
In acidic medium: H2O2 + 2H+ + 2e- → 2H2O
In basic medium :H2O2 + OH- + 2e- → 3OH--
As reducing agent
In acidic medium: H2O2 → 2H+ + O2 + 2e-
In basic medium : H2O2 + 2OH- → 2H2O + O2 + 2e-
2Fe2+ + H2O2 + 2H+ → 2Fe3+ + 2H2O
2MnO4- + 5H2O2 + 6H+ → 2Mn2+ + 8H2O + 5O2
Mn2+ + H2O2 → Mn+4 + 2OH-
2Fe3+ + H2O2 + 2OH- → 2Fe2+ + 2H2O + O2
The oxidising property of hydrogen peroxide is put to use in the restoration of old paintings, where the original white lead paint has been converted to black PbS by the H2S in the atmosphere. Hydrogen peroxide oxidises the black PbS into white PbSO4.
PbS(s) + 4H2O2 (aq) → PbSO4(s) + 4H2O
black white
c. Tests :
· It liberates iodine from potassium iodide in presence of ferrous sulphate
· Acidified solution of dichromate ion forms a deep blue colour with H2O2 due to the formation of CrO5. ,
Cr2O72- + 4H2O2 + 2H+ → 2CrO5 +5H2O
· With a solution of titanium oxide in conc.H2SO4, it gives orange colour due to the formation of pertitanic acid.
Ti4+ + H2O2 + 2H2O → H2TiO4
+ 4H+
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