Hydrogen and Its Compounds & S-Block Elements: Detailed Notes

Part 1: Hydrogen and Its Compounds

I. Position of Hydrogen in the Periodic Table Hydrogen, with its unique electronic configuration (1s^1), holds a special and somewhat ambiguous position in the periodic table. It exhibits properties that resemble both:

• Alkali Metals (Group 1): Like alkali metals, hydrogen has only one electron in its valence shell and readily loses it to form a unipositive ion (H^+). It combines with electronegative elements (like oxygen and halogens) to form compounds such as H2O, HCl, and also forms oxides and halides similar to those of alkali metals. It has an oxidation state of +1.
• Halogens (Group 17): Hydrogen requires only one electron to complete its duplet (achieve a stable noble gas configuration like Helium). It can gain an electron to form a uninegative ion (H^-, hydride ion), similar to how halogens form halide ions (X^-). It exists as a diatomic molecule (H2) and has an ionization energy comparable to some halogens. It can exhibit an oxidation state of -1 in metallic hydrides.

Despite these similarities, hydrogen’s exceptionally small size, very high ionization energy (much higher than alkali metals), and lack of d-orbitals set it apart. It doesn’t truly belong to either group, hence its unique placement.

II. Isotopes of Hydrogen Hydrogen has three naturally occurring isotopes, differing in the number of neutrons in their nucleus:

1. Protium (¹H): This is the most abundant isotope, making up about 99.985% of natural hydrogen. Its nucleus contains only one proton and no neutrons.
2. Deuterium (²H or D): Also known as heavy hydrogen, its nucleus contains one proton and one neutron. It is non-radioactive and is a stable isotope. Compounds formed with deuterium are often referred to as “deuterated” compounds. Deuterium oxide (D2O), or heavy water, is notably used as a moderator in nuclear reactors to slow down fast neutrons and as a tracer in chemical and biological reactions.
3. Tritium (³H or T): This isotope is radioactive, containing one proton and two neutrons in its nucleus. It is a beta (β) emitter, decaying with a relatively short half-life of 12.33 years. Tritium is used in self-powered lighting, as a radioactive tracer, and in nuclear fusion research.

III. Preparation of Dihydrogen (H2) Dihydrogen can be prepared using various laboratory and industrial methods:

1. Laboratory Methods:
   • From Metals and Acids: The most common laboratory method involves the reaction of active metals (like Zinc, Magnesium, Iron) with dilute non-oxidizing acids. Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g) Fe(s) + H2SO4(aq) → FeSO4(aq) + H2(g)
   • From Metals and Bases (Alkali/Alkaline Earth Metal Hydroxides): Certain amphoteric metals (like Zinc, Aluminum) react with strong alkalis to liberate hydrogen. 2NaOH(aq) + Zn(s) → Na2ZnO2(aq) + H2(g) (Sodium Zincate) 2Al(s) + 2NaOH(aq) + 6H2O(l) → 2NaAl(OH)4 + 3H2(g) (Sodium Tetrahydroxoaluminate(III))
2. Industrial Methods:
   • Electrolysis of Acidified Water: High purity dihydrogen (and dioxygen) can be produced by the electrolysis of acidified water (using a small amount of H2SO4 or NaOH to increase conductivity). 2H2O(l) –electrolysis–> 2H2(g) + O2(g)
   • From Brine Electrolysis (NaOH Production): Dihydrogen is obtained as a significant byproduct during the electrolysis of concentrated aqueous sodium chloride (brine) for the production of sodium hydroxide and chlorine. 2NaCl(aq) + 2H2O(l) –electrolysis–> 2NaOH(aq) + Cl2(g) + H2(g)
   • Steam Reforming of Hydrocarbons (Syngas Production): This is a major industrial method. Methane (natural gas) or other light hydrocarbons react with steam at high temperatures (around 1270K) in the presence of a nickel catalyst to produce a mixture of carbon monoxide and dihydrogen, known as ‘Syngas’ or ‘Water Gas’. CH4(g) + H2O(g) –Ni catalyst, 1270K–> CO(g) + 3H2(g) For coal: C(s) + H2O(g) –1270K–> CO(g) + H2(g)
   • Water-Gas Shift Reaction: To maximize the yield of hydrogen from syngas, the carbon monoxide component is further reacted with steam in the ‘water-gas shift reaction’ using an iron-chromium catalyst at 673K. This converts CO into more H2 and CO2. CO(g) + H2O(g) –Iron-chromium catalyst, 673K–> CO2(g) + H2(g) The carbon dioxide (CO2) produced is then selectively removed by scrubbing the gas mixture with a suitable solvent, often an aqueous solution of sodium arsenite.
   • From Coal Gasification: This process involves the conversion of coal into syngas. It is less environmentally friendly than steam reforming due to high CO2 emissions, but is a source of syngas.

IV. Properties of Dihydrogen (H2)

1. Physical Properties: Dihydrogen is a colorless, odorless, and tasteless gas. It is the lightest element known, making it buoyant. It is virtually insoluble in water. Dihydrogen has exceptionally high thermal conductivity, which has applications in cooling electrical generators. Its critical temperature is very low (33K), meaning it’s difficult to liquefy.
2. Chemical Properties: The H-H bond dissociation enthalpy is very high (435.88 kJ/mol), making dihydrogen relatively unreactive at room temperature. However, it reacts under specific conditions (high temperature, catalysts) due to its high bond strength.
   • Combustibility: Dihydrogen is highly flammable and burns readily in air (or oxygen) with a characteristic pale blue flame to produce water. This reaction is highly exothermic. 2H2(g) + O2(g) → 2H2O(l)
   • Reducing Nature: Dihydrogen is a good reducing agent, particularly for metal oxides of less reactive metals and non-metal compounds. Fe3O4(s) + 4H2(g) → 3Fe(s) + 4H2O(l) CuO(s) + H2(g) → Cu(s) + H2O(l)
   • Reaction with Halogens (X2): Dihydrogen reacts with halogens to form hydrogen halides (HX). The reactivity decreases down the group (F2 > Cl2 > Br2 > I2). The reaction with fluorine is explosive even at low temperatures, while with iodine it requires a catalyst and heat. H2(g) + X2(g) → 2HX(g) (where X = F, Cl, Br, I)
   • Reaction with Metals: Dihydrogen reacts with highly electropositive metals (especially s-block elements) at high temperatures to form ionic (saline) hydrides, where hydrogen exists as the hydride ion (H-). 2Na(s) + H2(g) → 2NaH(s) Ca(s) + H2(g) → CaH2(s)
   • Reaction with Organic Compounds (Hydrogenation): Dihydrogen is widely used in organic chemistry:
     • Hydrogenation of Unsaturated Hydrocarbons: Used to convert unsaturated compounds (containing C=C or C≡C bonds) into saturated ones. For example, the conversion of vegetable oils (liquid) to vanaspati ghee (solid fat) using a nickel (Ni), platinum (Pt), or palladium (Pd) catalyst.
     • Hydroformylation of Olefins (Oxo Process): Dihydrogen and carbon monoxide are added to olefins to produce aldehydes, which can then be reduced to alcohols.
     • Haber’s Process for Ammonia Synthesis: This is a cornerstone industrial process where nitrogen (N2) reacts with dihydrogen (H2) under high pressure and temperature, in the presence of an iron catalyst (with molybdenum as a promoter), to produce ammonia. N2(g) + 3H2(g) <=> 2NH3(g)

V. Hydrides Hydrides are binary compounds formed between hydrogen and other elements. They are broadly classified into three main types based on their chemical and physical properties:

1. Ionic (Saline) Hydrides:
   • Formation: Formed by the highly electropositive s-block elements (Group 1 and 2, specifically alkali metals and heavier alkaline earth metals like Ca, Sr, Ba). Beryllium (Be) and Magnesium (Mg) form covalent hydrides.
   • Nature: They are crystalline, non-volatile solids. In the solid state, they are non-conductors of electricity because the ions are fixed in the lattice. However, in the molten state or when dissolved in a suitable solvent, they conduct electricity due to the presence of mobile H- ions.
   • Reactivity: They are very reactive towards water and other protonic solvents, undergoing hydrolysis to produce dihydrogen gas. This strong reactivity makes them good reducing agents. NaH(s) + H2O(l) → NaOH(aq) + H2(g) CaH2(s) + 2H2O(l) → Ca(OH)2(aq) + 2H2(g) (Calcium hydride, “Hydrolith,” used for portable H2 generation)
2. Covalent (Molecular) Hydrides:
   • Formation: These are typically formed by elements of the p-block (Groups 13-17).
   • Nature: They are volatile compounds (gases or liquids at room temperature) with discrete molecules. The bonding is covalent.
   • Sub-classification based on electron count:
     • Electron-deficient Hydrides: Formed by Group 13 elements (e.g., Boron hydrides like B2H6 – diborane). These hydrides do not have enough electrons to form conventional two-center, two-electron (2c-2e) bonds for all atoms. They often feature multicenter bonding (e.g., 3-center, 2-electron bonds). They act as Lewis acids (electron acceptors).
     • Electron-precise Hydrides: Formed by Group 14 elements (e.g., CH4, SiH4). These hydrides have the exact number of electrons required to form standard 2c-2e covalent bonds. Their Lewis structures are complete.
     • Electron-rich Hydrides: Formed by Group 15, 16, and 17 elements (e.g., NH3, H2O, HF). These hydrides have one or more lone pairs of electrons on the central atom in addition to the bonding pairs. They can act as Lewis bases (electron donors) and are capable of forming intermolecular hydrogen bonds, significantly influencing their physical properties.
3. Metallic (Non-stoichiometric or Interstitial) Hydrides:
   • Formation: Primarily formed by elements of the d-block and f-block. Group 7, 8, and 9 elements typically do not form hydrides (this is known as the “Hydride Gap”).
   • Nature: These are often non-stoichiometric compounds, meaning their composition does not conform to simple whole-number ratios (e.g., LaH2.87, TiH1.7). They are formed by hydrogen atoms occupying the interstitial sites (gaps) within the metal lattice, without forming true chemical bonds.
   • Properties: Their properties vary. They are typically good conductors of heat and electricity. Some of them (e.g., Pd and Pt hydrides) can absorb large volumes of hydrogen, making them potential candidates for hydrogen storage. The absorbed hydrogen can be released upon heating.

VI. Water (H2O) Water is one of the most vital chemical compounds, essential for all known forms of life.

1. Structure: The water molecule has a bent or V-shape. The central oxygen atom undergoes sp3 hybridization, with two hybrid orbitals forming sigma bonds with two hydrogen atoms and the other two accommodating lone pairs of electrons. The presence of these two lone pairs leads to lone pair-lone pair (lp-lp) repulsion > lone pair-bond pair (lp-bp) repulsion > bond pair-bond pair (bp-bp) repulsion, which compresses the H-O-H bond angle from the ideal tetrahedral angle (109.5°) to approximately 104.5°. The molecule is highly polar due to the bent shape and the electronegativity difference between oxygen and hydrogen.
2. Hydrogen Bonding: Water exhibits extensive intermolecular hydrogen bonding. Each water molecule can form up to four hydrogen bonds with neighboring water molecules. This strong intermolecular force is responsible for many of water’s unusual physical properties compared to other hydrides of Group 16 elements (H2S, H2Se, H2Te).
3. Physical Properties:
   • High Melting and Boiling Points: Much higher than expected for a molecule of its size due to strong hydrogen bonding.
   • High Specific Heat Capacity: Water can absorb a large amount of heat without a significant rise in temperature, making it an excellent temperature regulator for living organisms and climates.
   • High Heat of Vaporization: Requires a large amount of energy to convert liquid water into vapor, important for evaporative cooling.
   • High Dielectric Constant (78.39 at 298 K): This high value allows water to effectively reduce the electrostatic forces of attraction between oppositely charged ions, making it an excellent universal solvent for many ionic and polar covalent compounds.
   • Anomalous Expansion (Ice Floats): Unlike most substances, water expands upon freezing. Liquid water has maximum density at 4°C. Below 4°C, it expands, and ice (0°C) is less dense than liquid water, which is why ice floats. This is crucial for aquatic life, as ice forms on the surface of water bodies, insulating the water below and preventing it from freezing completely. This open cage-like structure in ice is a result of hydrogen bonding.
4. Chemical Properties:
   • Amphoteric Nature: Water can act as both an acid (proton donor) and a base (proton acceptor), according to the Brønsted-Lowry theory.
     • As an acid: H2O(l) + NH3(aq) <=> OH-(aq) + NH4+(aq) (Water donates a proton to ammonia)
     • As a base: H2O(l) + H2S(aq) <=> H3O+(aq) + HS-(aq) (Water accepts a proton from H2S)
   • Redox Reactions: Water can be involved in redox reactions, acting as both an oxidizing and reducing agent.
     • With electropositive metals (oxidizing agent): 2Na + 2H2O → 2NaOH + H2 (Water is reduced to H2)
     • With strong oxidizing agents like F2 (reducing agent): 2F2 + 2H2O → 4HF + O2 (Water is oxidized to O2)
   • Hydrolysis Reactions: Water can react with many ionic and covalent compounds, breaking them down through hydrolysis.
     • With non-metal oxides: P4O10 + 6H2O → 4H3PO4 (forms phosphoric acid)
     • With chlorides: SiCl4 + 2H2O → SiO2 + 4HCl (forms silicic acid, precipitates SiO2)
     • With carbides: CaC2 + 2H2O → Ca(OH)2 + C2H2 (forms acetylene)
5. Hard and Soft Water:
   • Hard Water: Water containing dissolved mineral salts, primarily bicarbonate, sulfate, and chloride salts of calcium (Ca2+) and magnesium (Mg2+). Hard water does not easily form lather with soap because the Ca2+ and Mg2+ ions react with soap (sodium stearate, C17H35COONa) to form insoluble precipitates (scum).
     • Temporary Hardness: Caused by the presence of dissolved bicarbonates of calcium and magnesium (Ca(HCO3)2 and Mg(HCO3)2). It can be removed by simple methods.
       • Boiling: Heating decomposes the bicarbonates into insoluble carbonates, which precipitate out. Ca(HCO3)2 –heat–> CaCO3(s) + H2O(l) + CO2(g)
       • Clark’s Method: Adding a calculated amount of slaked lime (calcium hydroxide, Ca(OH)2). This precipitates both calcium carbonate and magnesium hydroxide. Ca(HCO3)2 + Ca(OH)2 → 2CaCO3(s) + 2H2O(l) Mg(HCO3)2 + 2Ca(OH)2 → 2CaCO3(s) + Mg(OH)2(s) + 2H2O(l)
     • Permanent Hardness: Caused by the presence of dissolved sulfates and chlorides of calcium and magnesium (CaSO4, MgSO4, CaCl2, MgCl2). It cannot be removed by boiling. Methods for removal include:
       • Washing Soda Method (Sodium Carbonate, Na2CO3): Sodium carbonate precipitates the Ca2+ and Mg2+ ions as insoluble carbonates. CaSO4 + Na2CO3 → CaCO3(s) + Na2SO4
       • Calgon Method (Sodium Hexametaphosphate, (NaPO3)6): Calgon forms soluble complex ions with Ca2+ and Mg2+, sequestering them from reacting with soap.
       • Ion-Exchange Method (Zeolite or Permutit Method, Synthetic Resins): This highly effective method involves passing hard water through a column containing ion-exchange resins. The resins (e.g., RNa, where R is a complex organic or inorganic anionic group) exchange their Na+ ions for the Ca2+ and Mg2+ ions in the water, making the water soft. The exhausted resin can be regenerated.

• Soft Water: Water that is free from dissolved Ca and Mg salts. It readily forms lather with soap.

VII. Hydrogen Peroxide (H2O2) Hydrogen peroxide is a powerful oxidizing agent and a widely used chemical.

1. Preparation:
   • From Barium Peroxide (BaO2): Historically, pure H2O2 was prepared by reacting hydrated barium peroxide with cold dilute sulfuric acid. BaO2.8H2O(s) + H2SO4(aq) → BaSO4(s) + H2O2(aq) + 8H2O(l)
   • Electrolytic Process: Industrially, H2O2 is prepared by the electrolysis of a cold, concentrated (50%) solution of sulfuric acid or an ammonium sulfate solution. Peroxodisulfuric acid (H2S2O8) is formed at the anode, which on hydrolysis yields H2O2. At anode: 2H2SO4 → H2S2O8 + 2H+ + 2e- (or 2SO4^2- → S2O8^2- + 2e-) H2S2O8 + 2H2O → 2H2SO4 + H2O2
   • Auto-oxidation of 2-ethylanthraquinol: This is the most common modern industrial method. 2-ethylanthraquinol is catalytically oxidized by air to 2-ethylanthraquinone and H2O2. The 2-ethylanthraquinone is then reduced back to 2-ethylanthraquinol, recycling the organic compound.
2. Structure: Hydrogen peroxide has a non-planar, open book-like structure. The two O-H bonds are in different planes.
   • In the gas phase, the H-O-O-H dihedral angle is approximately 111.5°.
   • In the solid phase (crystalline), the dihedral angle is approximately 90.2°. This non-planar structure is due to repulsion between the two lone pairs on the oxygen atoms.
3. Physical Properties: H2O2 is a colorless, syrupy liquid. It has a higher density and viscosity than water. It is miscible with water in all proportions due to its ability to form hydrogen bonds with water. Pure H2O2 freezes at -0.41°C and boils at 150.2°C. It is a powerful oxidizing agent. It decomposes slowly in the presence of light, metals, and bases, which catalyze its decomposition. For this reason, it is stored in dark, wax-lined plastic bottles or in glass bottles with a little urea (which acts as a stabilizer).
4. Chemical Properties: Hydrogen peroxide is a versatile compound, capable of acting as both an oxidizing and a reducing agent, depending on the reaction conditions and the nature of the reacting species.
   • Decomposition: H2O2 is thermodynamically unstable and disproportionates into water and oxygen. This decomposition is accelerated by light, rough surfaces, metal ions (e.g., Fe2+, Mn2+), and bases. 2H2O2(l) → 2H2O(l) + O2(g)
   • Oxidizing Agent: H2O2 acts as an oxidizing agent in both acidic and basic media.
     • In acidic medium: 2Fe2+ + H2O2 + 2H+ → 2Fe3+ + 2H2O (Oxidizes Fe2+ to Fe3+) PbS(s) + 4H2O2(aq) → PbSO4(s) + 4H2O(l) (Oxidizes black lead sulphide to white lead sulphate)
     • In basic medium: 2Fe2+ + H2O2 → 2Fe3+ + 2OH- (Oxidizes Fe2+ to Fe3+) Mn2+ + H2O2 → Mn4+ (as MnO2)
   • Reducing Agent: H2O2 can also act as a reducing agent, especially when reacting with strong oxidizing agents.
     • In acidic medium: 2MnO4- + 5H2O2 + 6H+ → 2Mn2+ + 5O2 + 8H2O (Reduces permanganate) Cl2 + H2O2 → 2HCl + O2 (Reduces chlorine)
     • In basic medium: 2MnO4- + 3H2O2 → 2MnO2 + 3O2 + 2H2O + 2OH- (Reduces permanganate to MnO2)
   • Bleaching Action: H2O2 is an effective bleaching agent for delicate materials like textiles, paper pulp, hair, and wool. Its bleaching action is due to the nascent oxygen released during its decomposition. H2O2 → H2O + [O] (nascent oxygen) Colored matter + [O] → Colorless matter
5. Uses: Hydrogen peroxide has diverse applications: as an antiseptic (e.g., for wounds under the name “perhydrol”), a bleaching agent, an oxidizing agent in various chemical syntheses, a rocket fuel, and in environmental applications (e.g., treatment of industrial effluents).

Part 2: S-Block Elements

The s-block elements consist of Group 1 (Alkali Metals) and Group 2 (Alkaline Earth Metals). Their characteristic properties are due to their outermost s-orbital containing one (Group 1) or two (Group 2) valence electrons.

I. Group 1: Alkali Metals (Li, Na, K, Rb, Cs, Fr) These are highly reactive metals due to their tendency to readily lose their single valence electron. Francium (Fr) is radioactive.

• Electronic Configuration: Their general electronic configuration is [Noble gas]ns^1. They readily lose this single electron to form stable unipositive ions (M^+).
• Atomic and Ionic Radii: Both atomic and ionic radii increase consistently down the group. This is due to the addition of new electron shells with increasing atomic number. The ionic radius (M^+) is always smaller than the corresponding atomic radius (M) because of the loss of the outermost electron and increased effective nuclear charge.
• Ionization Enthalpy (IE): Alkali metals have very low first ionization enthalpies, indicating that it is very easy to remove their outermost electron. This value decreases steadily down the group (Li > Na > K > Rb > Cs) as atomic size increases and the valence electron is further from the nucleus.
• Hydration Enthalpy: The hydration enthalpy of alkali metal ions decreases down the group (Li+ > Na+ > K+ > Rb+ > Cs+). This is because hydration enthalpy is inversely proportional to the ionic size; smaller ions are more intensely hydrated (surrounded by water molecules) due to a higher charge density. Lithium ion (Li+) has the highest hydration enthalpy.
• Density: Generally increases down the group. However, there is an anomaly: Potassium (K) is less dense than Sodium (Na). This is due to the unexpectedly large increase in atomic volume from Na to K, which outweighs the increase in atomic mass.
• Melting and Boiling Points: These metals have low melting and boiling points, which decrease down the group (Li > Na > K > Rb > Cs). This is because of their weak metallic bonding, as they only contribute one electron per atom to the metallic bond. As atomic size increases, the strength of this metallic bond further weakens.
• Flame Coloration: Alkali metals and their salts impart characteristic and persistent colors to a non-luminous flame (e.g., Bunsen burner flame). This is due to the excitation of their outermost electron to a higher energy level by the heat of the flame, followed by the emission of light at specific wavelengths as the electron returns to its ground state.
  • Lithium (Li): Crimson red
  • Sodium (Na): Golden yellow (most intense and commonly used for identification)
  • Potassium (K): Lilac (pale violet)
  • Rubidium (Rb): Red-violet
  • Cesium (Cs): Blue
• Reducing Nature: Alkali metals are powerful reducing agents because of their strong tendency to lose electrons. Lithium (Li) is the strongest reducing agent in aqueous solution. Although it has the highest ionization energy, its exceptionally high hydration enthalpy (due to its small size) makes the overall standard electrode potential highly negative, favoring reduction.
• Chemical Properties:
  • Reaction with Air/Oxygen: Alkali metals tarnish in dry air due to the formation of oxides. They react vigorously with oxygen.
    • Lithium (Li) primarily forms a normal oxide (monoxide): 4Li + O2 → 2Li2O
    • Sodium (Na) forms a peroxide: 2Na + O2 → Na2O2
    • Potassium (K), Rubidium (Rb), and Cesium (Cs) form superoxides: M + O2 → MO2 (where M = K, Rb, Cs). Superoxides contain the O2^- ion.
  • Reaction with Water: Alkali metals react vigorously with water, producing metal hydroxides and dihydrogen gas. The reactivity increases down the group. Lithium reacts relatively slowly (due to strong bonding in its small size and high lattice energy), while Cesium reacts explosively. 2M(s) + 2H2O(l) → 2MOH(aq) + H2(g) + Heat
  • Reaction with Hydrogen: They react with dihydrogen at about 673K (300°C) to form ionic hydrides (MH), which are crystalline solids containing the H^- ion. 2M(s) + H2(g) → 2MH(s)
  • Reaction with Halogens: They react readily with halogens to form ionic halides (MX). The reactivity increases down the group. 2M(s) + X2(g) → 2MX(s)
  • Solutions in Liquid Ammonia: Alkali metals dissolve in liquid ammonia, forming deep blue, electrically conducting solutions. This blue color is due to the presence of ammoniated electrons, which absorb energy in the visible region. These solutions are strongly reducing and paramagnetic. At higher concentrations, the blue color changes to bronze, and the solutions become diamagnetic due to the pairing of ammoniated electrons. Upon standing, these solutions slowly decompose to form amides and liberate dihydrogen. M(s) + (x+y)NH3(l) → [M(NH3)x]+(am) + [e(NH3)y]-(am) (blue, conducting) 2M(s) + 2NH3(l) → 2MNH2(am) + H2(g) (slow decomposition to amide)

Anomalous Properties of Lithium: Lithium, the first element of Group 1, shows distinct anomalous properties compared to the rest of the alkali metals, primarily due to its:

• Extremely small size.
• High ionization enthalpy.
• High polarizing power (ability to distort the electron cloud of an anion). These properties lead to:
• Formation of monoxide (Li2O), not peroxide/superoxide.
• Less vigorous reaction with water compared to other alkali metals.
• Formation of a very stable nitride (Li3N) by direct reaction with nitrogen (unique among alkali metals).
• Formation of more covalent compounds (e.g., LiCl is more covalent and soluble in organic solvents like ethanol, unlike other alkali metal chlorides).
• It does not form ethynides on reaction with acetylene.

Diagonal Relationship with Magnesium (Mg): Lithium exhibits a diagonal relationship with Magnesium (Group 2, period 3). This is because they have similar charge-to-size ratios, leading to comparable polarizing powers and thus similar properties.

• Both form stable nitrides (Li3N, Mg3N2) by direct combination with nitrogen.
• Both have covalent chlorides (LiCl, MgCl2) that are deliquescent and soluble in ethanol.
• Both are harder and have higher melting points than other elements in their respective groups.
• Their hydroxides (LiOH, Mg(OH)2) are weak bases and sparingly soluble.
• Both form basic oxides.
• Both form complexes (e.g., Li[AlF6], Mg[AlF6]).
• Both carbonates (Li2CO3, MgCO3) decompose easily on heating.
• Both metals react slowly with water.

Important Compounds of Sodium:

• Sodium Carbonate (Washing Soda, Na2CO3.10H2O): Industrially prepared by the Solvay process (Ammonia-Soda process). It is used extensively in the manufacture of glass, soap, paper, textiles, and for softening hard water.
• Sodium Chloride (NaCl): Common table salt, obtained from seawater or rock salt deposits. It’s a crucial raw material for the production of NaOH, Na2CO3, NaHCO3, and other sodium compounds.
• Sodium Hydroxide (Caustic Soda, NaOH): Manufactured by the electrolysis of concentrated aqueous NaCl solution (brine) using the Castner-Kellner cell. It’s a strong base widely used in soap and detergent manufacturing, paper industry, textile industry, and refining of petroleum.
• Sodium Bicarbonate (Baking Soda, NaHCO3): An intermediate in the Solvay process, it is also found naturally. Used in baking powders (reacts with acids to release CO2), as an antacid (neutralizes stomach acid), and in fire extinguishers.

II. Group 2: Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra) These elements are also quite reactive, but generally less so than alkali metals. Radium (Ra) is radioactive.

• Electronic Configuration: Their general electronic configuration is [Noble gas]ns^2. They readily lose both valence electrons to form stable dipositive ions (M^2+).
• Atomic and Ionic Radii: Both atomic and ionic radii are smaller than their corresponding alkali metals in the same period (due to higher nuclear charge) but increase down the group due to the addition of new electron shells.
• Ionization Enthalpy (IE): Their first ionization enthalpies are higher than those of alkali metals (due to smaller size and greater nuclear charge), but they still have relatively low ionization energies. The IE decreases down the group. The second ionization enthalpy is always higher than the first because it involves removing an electron from a positively charged ion.
• Hydration Enthalpy: The hydration enthalpy of alkaline earth metal ions (M^2+) decreases down the group (Be2+ > Mg2+ > Ca2+ > Sr2+ > Ba2+). Be2+ has the highest hydration enthalpy due to its smallest size and highest charge density, similar to Li+.
• Density: Generally increases down the group, but there are some minor irregularities (e.g., Ca is less dense than Mg).
• Melting and Boiling Points: They generally have higher melting and boiling points than alkali metals of the same period because they contribute two valence electrons to the metallic bond, leading to stronger metallic bonding. The melting and boiling points generally decrease down the group.
• Flame Coloration: Like alkali metals, heavier alkaline earth metals impart characteristic colors to a non-luminous flame. Beryllium (Be) and Magnesium (Mg) do not show flame coloration because their electrons are more tightly held (due to smaller size and higher ionization energy), and the energy of the flame is insufficient to excite them to higher energy levels.
  • Calcium (Ca): Brick red
  • Strontium (Sr): Crimson red
  • Barium (Ba): Apple green
• Reducing Nature: Alkaline earth metals are strong reducing agents, although they are less powerful than alkali metals (due to higher ionization energies). Their reducing power generally increases down the group as ionization enthalpy decreases. Beryllium is the weakest reducing agent among them.
• Chemical Properties:
  • Reaction with Air/Oxygen: They react with air to form oxides (MO). Beryllium and Magnesium form a protective oxide layer that prevents further reaction. Barium forms a peroxide (BaO2). Magnesium and Beryllium can also react with nitrogen at high temperatures to form nitrides (Mg3N2, Be3N2). 2M(s) + O2(g) → 2MO(s) 2M(s) + N2(g) → M3N2(s) (where M=Be, Mg)
  • Reaction with Water: They are less reactive with water compared to alkali metals. Beryllium does not react with water or steam. Magnesium reacts slowly with cold water but readily with steam. Calcium, Strontium, and Barium react increasingly vigorously with cold water. M(s) + 2H2O(l) → M(OH)2(aq) + H2(g)
  • Reaction with Halogens: They react with halogens at elevated temperatures to form ionic halides (MX2). However, Beryllium chloride (BeCl2) is predominantly covalent due to the high polarizing power of the small Be2+ ion. M(s) + X2(g) → MX2(s)
  • Reaction with Hydrogen: They form ionic hydrides (MH2) at elevated temperatures. M(s) + H2(g) → MH2(s)
  • Solutions in Liquid Ammonia: Like alkali metals, alkaline earth metals also dissolve in liquid ammonia to form deep blue solutions (containing ammoniated ions and electrons), but these solutions are less stable and decompose more readily than those of alkali metals.

Anomalous Properties of Beryllium: Beryllium, the first element of Group 2, displays anomalous behavior due to its:

• Exceptionally small size.
• High ionization enthalpy.
• High polarizing power. These properties lead to:
• Formation of predominantly covalent compounds (e.g., BeCl2, BeO).
• Does not react with water or steam, unlike other alkaline earth metals.
• Its oxide (BeO) and hydroxide (Be(OH)2) are amphoteric, reacting with both acids and bases.
• Forms stable carbides (Be2C) by direct combination with carbon.
• Forms complexes more readily than other alkaline earth metals.

Diagonal Relationship with Aluminum (Al): Beryllium exhibits a diagonal relationship with Aluminum (Group 13, period 3). This is due to their similar charge-to-size ratios, resulting in comparable properties.

• Both form stable covalent halides (BeCl2, AlCl3) that exist as dimeric bridged structures (Be2Cl4, Al2Cl6) in the vapor phase and are strong Lewis acids.
• Both have amphoteric oxides (BeO, Al2O3) and hydroxides (Be(OH)2, Al(OH)3) that react with acids and strong bases.
• Both do not react with water in bulk, due to the formation of a protective oxide layer.
• Both form complexes (e.g., [BeF4]2-, [AlF6]3-).
• Both dissolve in strong alkalis to form soluble complex ions (e.g., [Be(OH)4]2-, [Al(OH)4]-).

Important Compounds of Calcium:

• Calcium Oxide (Quicklime, CaO): Produced on an industrial scale by heating limestone (CaCO3) in a lime kiln at around 1070-1270K. It is a basic oxide. Used extensively in cement manufacturing, as a basic flux in metallurgy, in glass and sugar industries, and for purifying water. CaCO3(s) –heat–> CaO(s) + CO2(g)
• Calcium Hydroxide (Slaked Lime, Ca(OH)2): Formed when quicklime reacts with water (slaking of lime). It is a white amorphous powder. Its aqueous solution is called limewater, and its suspension in water is called milk of lime. Used in whitewashing, neutralization of acidic soils, and in the preparation of mortar. CaO(s) + H2O(l) → Ca(OH)2(s)
• Calcium Carbonate (Limestone, CaCO3): Found naturally in various forms like limestone, marble, and chalk. It is a primary component of rocks. Used as a building material, in the manufacture of cement and quicklime, and as a filler.
• Calcium Sulphate (Gypsum, CaSO4.2H2O): A common mineral. Used in cement manufacturing to retard its setting time. When heated carefully to 393K (120°C), it loses three-quarters of its water of crystallization to form Plaster of Paris. CaSO4.2H2O(s) –393K–> CaSO4.1/2H2O(s) + 3/2 H2O(g)
  • Plaster of Paris (PoP, CaSO4.1/2H2O): A white powder that, when mixed with water, sets into a hard mass. Used in making casts for fractures, dental impressions, statues, and decorative materials. If gypsum is heated above 393K, it forms anhydrous calcium sulfate (CaSO4), known as “dead burnt plaster,” which loses its setting properties.

70 Important MCQs on Hydrogen and Its Compounds & S-Block Elements

Part A: Hydrogen and Its Compounds (35 Questions)

1. Which isotope of hydrogen is radioactive? (a) Protium (b) Deuterium (c) Tritium (d) All of the above
2. The maximum number of isotopes possible for hydrogen is: (a) 1 (b) 2 (c) 3 (d) 4
3. Water gas is a mixture of: (a) CO + N2 (b) CO + H2 (c) CO2 + H2 (d) CO2 + N2
4. The process used to remove temporary hardness of water is: (a) Calgon method (b) Ion-exchange method (c) Boiling (d) Washing soda method
5. Hydrogen peroxide (H2O2) has an open book-like structure. The dihedral angle in the gas phase is approximately: (a) 90.2° (b) 104.5° (c) 111.5° (d) 120°
6. Which of the following acts as a component for removing CO2 in the water gas shift reaction? (a) Iron-chromium catalyst (b) Sodium arsenite solution (c) Steam (d) Carbon monoxide
7. Which of the following hydrides is electron-deficient? (a) CH4 (b) NH3 (c) B2H6 (d) H2O
8. Heavy water (D2O) is used in nuclear reactors as a: (a) Fuel (b) Coolant (c) Moderator (d) Control rod
9. Which of the following is an example of an interstitial hydride? (a) NaH (b) H2O (c) TiH1.7 (d) NH3
10. The reaction of an ionic hydride with water produces: (a) Oxygen gas (b) Hydrogen gas (c) Water gas (d) No gas
11. What is the approximate bond angle in a water molecule? (a) 180° (b) 120° (c) 109.5° (d) 104.5°
12. Hydrogen shows properties similar to both alkali metals and halogens. Which property is unique to hydrogen and makes its position in periodic table ambiguous? (a) Forms H+ ion (b) Forms H- ion (c) Exists as diatomic molecule (d) Small size and high ionization energy
13. The industrial preparation of dihydrogen involves the reaction of steam with hydrocarbons at high temperature in the presence of a catalyst. This process is called: (a) Electrolysis (b) Steam reforming (c) Haber’s process (d) Water gas shift reaction
14. Hydrogen peroxide is stored in wax-lined plastic bottles because: (a) It reacts with glass (b) It is sensitive to light (c) It is highly corrosive (d) Both (a) and (b)
15. In the Haber’s process for ammonia synthesis, dihydrogen is reacted with: (a) Oxygen (b) Nitrogen (c) Carbon monoxide (d) Water
16. Which of the following is not a primary use of hydrogen peroxide? (a) Bleaching agent (b) Antiseptic (c) Rocket fuel (d) As a strong reducing agent for non-oxidizing acids
17. Permanent hardness of water cannot be removed by: (a) Washing soda method (b) Ion-exchange method (c) Calgon method (d) Boiling
18. Which type of hydride is generally non-stoichiometric? (a) Ionic hydrides (b) Covalent hydrides (c) Metallic hydrides (d) All of the above
19. What happens when metallic sodium reacts with hydrogen gas? (a) Sodium hydroxide is formed (b) Sodium hydride is formed (c) Sodium carbonate is formed (d) No reaction occurs
20. The common name for the mixture of CO and H2 produced from coal or hydrocarbons is: (a) Producer gas (b) Syngas (c) Natural gas (d) Biogas
21. Water is an amphoteric substance, meaning it can act as: (a) Only an acid (b) Only a base (c) Both an acid and a base (d) Neither an acid nor a base
22. The anomalous behavior of water (ice floating on water) is due to: (a) High specific heat (b) High dielectric constant (c) Hydrogen bonding (d) Low heat of vaporization
23. Which of the following elements does not form a hydride (Hydride Gap)? (a) Scandium (b) Titanium (c) Manganese (d) Palladium
24. Hydrogen peroxide can act as: (a) Only an oxidizing agent (b) Only a reducing agent (c) Both an oxidizing and reducing agent (d) Neither an oxidizing nor a reducing agent
25. The process of converting vegetable oils to vanaspati ghee using dihydrogen is called: (a) Saponification (b) Esterification (c) Hydrogenation (d) Fermentation
26. The half-life of Tritium is approximately: (a) 5 years (b) 12.33 years (c) 25 years (d) 50 years
27. Which of the following is an electron-rich hydride? (a) BCl3 (b) SiH4 (c) HF (d) C2H6
28. The highest percentage of Protium in natural hydrogen is due to: (a) Its radioactivity (b) Its stability (c) Its higher atomic mass (d) Its lower atomic mass
29. The method used to remove CO2 from syngas in the water-gas shift reaction is by scrubbing with: (a) NaOH solution (b) H2SO4 solution (c) Sodium arsenite solution (d) Water
30. In the electrolytic preparation of H2O2, which acid is typically electrolysed at high concentration? (a) HCl (b) HNO3 (c) H2SO4 (d) H3PO4
31. Which of the following is not a property of ionic hydrides? (a) Crystalline solids (b) Non-volatile (c) Good conductors in solid state (d) React violently with water
32. The process of hydrogenation of vegetable oils requires a catalyst, typically: (a) Iron (b) Platinum (c) Nickel (d) Vanadium
33. Temporary hardness of water is caused by the presence of: (a) Sulfates of Ca and Mg (b) Chlorides of Ca and Mg (c) Bicarbonates of Ca and Mg (d) Carbonates of Na and K
34. Which property of hydrogen peroxide makes it useful as a bleaching agent? (a) Its strong oxidizing nature (b) Its strong reducing nature (c) Its acidic nature (d) Its basic nature
35. The reaction of dihydrogen with oxygen to form water is: (a) Endothermic (b) Exothermic (c) Reversible (d) Photochemical

Part B: S-Block Elements (35 Questions)

36. Which alkali metal has the highest hydration enthalpy? (a) Lithium (b) Sodium (c) Potassium (d) Cesium
37. Which of the following alkaline earth metals does not impart color to the flame? (a) Calcium (b) Strontium (c) Barium (d) Magnesium
38. The general electronic configuration of alkali metals is: (a) ns1 (b) ns2 (c) ns2np1 (d) (n-1)d1ns2
39. Which of the following compounds is prepared by the Solvay process? (a) NaOH (b) NaHCO3 (c) Na2CO3 (d) NaCl
40. Anomalous behavior of Lithium is due to: (a) Its small size and high polarizing power (b) Its high electronegativity (c) Its tendency to form ionic bonds (d) Its liquid state at room temperature
41. Which of the following is correct regarding the density of alkali metals down the group? (a) Increases regularly (b) Decreases regularly (c) Increases, but Potassium is lighter than Sodium (d) Decreases, but Sodium is lighter than Potassium
42. Which of the following is an amphoteric oxide? (a) Li2O (b) MgO (c) BeO (d) CaO
43. Which pair of elements shows a diagonal relationship? (a) Li and Na (b) Be and Mg (c) Li and Mg (d) Na and K
44. In the Castner-Kellner process, what is the product formed at the mercury cathode? (a) Chlorine gas (b) Hydrogen gas (c) Sodium amalgam (d) Sodium hydroxide
45. The reaction of alkali metals with liquid ammonia produces a deep blue solution. This color is due to: (a) Ammoniated metal ions (b) Ammoniated electrons (c) Formation of amides (d) Presence of free electrons
46. When gypsum (CaSO4.2H2O) is heated to 393K, it forms: (a) Dead burnt plaster (b) Quicklime (c) Plaster of Paris (d) Calcium hydroxide
47. Which alkali metal forms a peroxide when reacted with oxygen? (a) Lithium (b) Sodium (c) Potassium (d) Rubidium
48. Which of the following is not a property of alkaline earth metals? (a) Strong reducing agents (b) Form M2+ ions (c) Have higher ionization enthalpies than alkali metals (d) Generally form highly soluble hydroxides
49. Which of the following is a constituent of Duralumin alloy? (a) Iron (b) Copper (c) Lead (d) Tin
50. Quicklime is chemically: (a) Calcium carbonate (b) Calcium hydroxide (c) Calcium oxide (d) Calcium sulfate
51. The melting and boiling points of alkali metals decrease down the group. This is because: (a) Atomic size increases (b) Metallic bonding weakens (c) Ionization enthalpy decreases (d) Density increases
52. Which of the following nitrates of alkaline earth metals decomposes on heating to give NO2 and O2? (a) Be(NO3)2 (b) Mg(NO3)2 (c) Ca(NO3)2 (d) All of the above
53. Alkali metals typically exhibit an oxidation state of: (a) +1 (b) +2 (c) +3 (d) -1
54. Which of the following is true about the basicity of hydroxides of alkaline earth metals down the group? (a) Increases (b) Decreases (c) Remains same (d) First increases then decreases
55. Which of the following is used in the manufacturing of cement? (a) Gypsum (b) Quicklime (c) Limestone (d) All of the above
56. The solubility of alkaline earth metal sulfates decreases down the group. This is due to: (a) Decrease in lattice enthalpy (b) Decrease in hydration enthalpy (c) Increase in lattice enthalpy (d) Increase in hydration enthalpy
57. When hydrated lithium chloride (LiCl.2H2O) is heated, it forms: (a) LiCl only (b) Li2O (c) LiOH (d) Li + Cl2
58. Which of the following alkaline earth metal salts is covalent? (a) MgCl2 (b) CaCl2 (c) BeCl2 (d) SrCl2
59. The reducing character of alkali metals in aqueous solution follows the order: (a) Li > Na > K > Rb > Cs (b) Cs > Rb > K > Na > Li (c) Na > K > Rb > Cs > Li (d) Li < Na < K < Rb < Cs
60. Slaked lime is chemically: (a) CaO (b) CaCO3 (c) Ca(OH)2 (d) CaSO4
61. Which of the following alkaline earth metal carbonates is most thermally stable? (a) BeCO3 (b) MgCO3 (c) CaCO3 (d) BaCO3
62. The biological importance of sodium ions is related to: (a) Muscle contraction (b) Nerve impulse transmission (c) Maintaining fluid balance (d) All of the above
63. Which of the following properties is characteristic of both Lithium and Magnesium? (a) Form stable nitrides (b) Form covalent chlorides soluble in ethanol (c) Hydroxides are weak bases (d) All of the above
64. Which of the following elements forms a superoxide on reaction with oxygen? (a) Na (b) K (c) Mg (d) Li
65. The alkaline earth metal that does not react with water even at high temperatures is: (a) Magnesium (b) Calcium (c) Beryllium (d) Strontium
66. Plaster of Paris is chemically: (b) Calcium sulfate hemihydrate (a) Calcium sulfate dihydrate (c) Calcium oxide (d) Calcium carbonate
67. Which of the following statements about the solubility of alkali metal halides (LiF, NaCl, KBr, CsI) in water is generally true? (a) Increases down the group (b) Decreases down the group (c) Follows no regular trend (d) LiF is most soluble
68. The industrial preparation of sodium hydroxide is by the: (a) Solvay process (b) Down’s process (c) Castner-Kellner process (d) Hall-Héroult process
69. Which of the following compounds of alkaline earth metals is used as an antacid? (b) Magnesium hydroxide (a) Calcium chloride (c) Barium sulfate (d) Strontium carbonate
70. The flame color produced by Barium is: (c) Apple green (a) Crimson red (b) Golden yellow (d) Lilac

Answer Key and Explanations for Hydrogen and Its Compounds & S-Block Elements MCQs

Part A: Hydrogen and Its Compounds

1. Which isotope of hydrogen is radioactive? (c) Tritium
   • Explanation: Tritium (³H) is the only radioactive isotope of hydrogen, undergoing beta decay.
2. The maximum number of isotopes possible for hydrogen is: (c) 3
   • Explanation: Hydrogen has three known isotopes: Protium (¹H), Deuterium (²H or D), and Tritium (³H or T).
3. Water gas is a mixture of: (b) CO + H2
   • Explanation: Water gas is produced by passing steam over hot coke or through hydrocarbons (steam reforming). It’s also known as syngas.
4. The process used to remove temporary hardness of water is: (c) Boiling
   • Explanation: Temporary hardness, caused by soluble bicarbonates of Ca and Mg, can be removed by boiling, which precipitates the insoluble carbonates.
5. Hydrogen peroxide (H2O2) has an open book-like structure. The dihedral angle in the gas phase is approximately: (c) 111.5°
   • Explanation: H2O2 has a non-planar structure. The dihedral angle between the two H-O planes is 111.5° in the gas phase and 90.2° in the solid phase.
6. Which of the following acts as a component for removing CO2 in the water gas shift reaction? (b) Sodium arsenite solution
   • Explanation: In the water-gas shift reaction, CO is converted to CO2 and H2. The CO2 produced is then removed by scrubbing with sodium arsenite solution.
7. Which of the following hydrides is electron-deficient? (c) B2H6
   • Explanation: Diborane (B2H6) is an electron-deficient compound. Boron hydrides are known for having insufficient valence electrons for conventional two-center, two-electron bonds.
8. **Heavy water (D2O) is used in nuclear reactors as a: ** (c) Moderator
   • Explanation: Heavy water (D2O) is used as a moderator in nuclear reactors to slow down fast neutrons, making them more effective in causing further fission.
9. Which of the following is an example of an interstitial hydride? (c) TiH1.7
   • Explanation: Metallic (or interstitial) hydrides are typically formed by d-block and f-block elements and are often non-stoichiometric, like TiH1.7.
10. The reaction of an ionic hydride with water produces: (b) Hydrogen gas
    • Explanation: Ionic hydrides (e.g., NaH, CaH2) are strong reducing agents and react violently with water to produce H2 gas and a metal hydroxide. For example, NaH + H2O → NaOH + H2.
11. What is the approximate bond angle in a water molecule? (d) 104.5°
    • Explanation: The water molecule has a bent shape due to two lone pairs on the central oxygen atom. These lone pairs repel the bonding pairs, reducing the bond angle from the ideal tetrahedral angle of 109.5° to approximately 104.5°.
12. Hydrogen shows properties similar to both alkali metals and halogens. Which property is unique to hydrogen and makes its position in periodic table ambiguous? (d) Small size and high ionization energy
    • Explanation: While hydrogen shares some characteristics with both alkali metals and halogens, its extremely small size and exceptionally high ionization energy differentiate it significantly from both groups, making its unique position a topic of discussion.
13. The industrial preparation of dihydrogen involves the reaction of steam with hydrocarbons at high temperature in the presence of a catalyst. This process is called: (b) Steam reforming
    • Explanation: Steam reforming of methane (CH4) is a major industrial method for producing syngas (CO + H2).
14. Hydrogen peroxide is stored in wax-lined plastic bottles because: (d) Both (a) and (b)
    • Explanation: H2O2 decomposes in the presence of light and also reacts with rough glass surfaces (alkali in glass catalyzes decomposition). Therefore, it’s stored in dark, wax-lined plastic bottles.
15. In the Haber’s process for ammonia synthesis, dihydrogen is reacted with: (b) Nitrogen
    • Explanation: The Haber-Bosch process synthesizes ammonia (NH3) from nitrogen (N2) and hydrogen (H2). N2(g) + 3H2(g) <=> 2NH3(g).
16. Which of the following is not a primary use of hydrogen peroxide? (d) As a strong reducing agent for non-oxidizing acids
    • Explanation: Hydrogen peroxide is generally an oxidizing agent. While it can act as a reducing agent in some reactions (e.g., with strong oxidizing agents like MnO4-), it is not primarily used as a strong reducing agent for non-oxidizing acids. Its main applications are as an oxidizing agent, bleach, and antiseptic.
17. Permanent hardness of water cannot be removed by: (d) Boiling
    • Explanation: Permanent hardness is caused by dissolved sulfates and chlorides of calcium and magnesium, which do not precipitate upon boiling.
18. Which type of hydride is generally non-stoichiometric? (c) Metallic hydrides
    • Explanation: Metallic (or interstitial) hydrides often have compositions that are not simple whole-number ratios, making them non-stoichiometric (e.g., TiH1.7, LaH2.87).
19. What happens when metallic sodium reacts with hydrogen gas? (b) Sodium hydride is formed
    • Explanation: Alkali metals react with hydrogen to form ionic (saline) hydrides, e.g., 2Na(s) + H2(g) → 2NaH(s).
20. The common name for the mixture of CO and H2 produced from coal or hydrocarbons is: (b) Syngas
    • Explanation: Syngas (synthesis gas) is a crucial industrial feedstock, primarily a mixture of carbon monoxide and hydrogen. It’s also known as water gas.
21. Water is an amphoteric substance, meaning it can act as: (c) Both an acid and a base
    • Explanation: Water can donate a proton (act as an acid) or accept a proton (act as a base), depending on the reacting species. E.g., H2O + NH3 → OH- + NH4+ (water is acid); H2O + H2S → H3O+ + HS- (water is base).
22. The anomalous behavior of water (ice floating on water) is due to: (c) Hydrogen bonding
    • Explanation: The extensive hydrogen bonding in water leads to an open, cage-like structure in ice, which makes ice less dense than liquid water, causing it to float.
23. Which of the following elements does not form a hydride (Hydride Gap)? (c) Manganese
    • Explanation: Elements of Group 7, 8, and 9 (e.g., Mn, Fe, Co, Ni) in the d-block do not form hydrides. This is known as the “hydride gap.”
24. Hydrogen peroxide can act as: (c) Both an oxidizing and reducing agent
    • Explanation: H2O2 is a versatile redox agent. It can oxidize substances (e.g., Fe2+ to Fe3+) and reduce others (e.g., MnO4- to Mn2+ or MnO2).
25. The process of converting vegetable oils to vanaspati ghee using dihydrogen is called: (c) Hydrogenation
    • Explanation: Hydrogenation is the addition of hydrogen to unsaturated compounds (like vegetable oils) to make them saturated (like vanaspati ghee), typically using a nickel catalyst.
26. The half-life of Tritium is approximately: (b) 12.33 years
    • Explanation: Tritium is a beta emitter with a half-life of 12.33 years.
27. Which of the following is an electron-rich hydride? (c) HF
    • Explanation: Electron-rich hydrides (from Groups 15, 16, 17) have lone pairs of electrons on the central atom (e.g., NH3, H2O, HF). BCl3 is electron-deficient, SiH4 and C2H6 are electron-precise.
28. The highest percentage of Protium in natural hydrogen is due to: (d) Its lower atomic mass
    • Explanation: Protium is the most abundant isotope because it is the simplest and lightest form of hydrogen, making it energetically favorable and thus most common.
29. The method used to remove CO2 from syngas in the water-gas shift reaction is by scrubbing with: (c) Sodium arsenite solution
    • Explanation: After the water-gas shift reaction, CO2 is selectively removed from the mixture by scrubbing with a solution of sodium arsenite.
30. In the electrolytic preparation of H2O2, which acid is typically electrolysed at high concentration? (c) H2SO4
    • Explanation: Industrially, H2O2 can be prepared by the electrolysis of an approximately 50% solution of H2SO4 or a solution of ammonium sulfate.
31. Which of the following is not a property of ionic hydrides? (c) Good conductors in solid state
    • Explanation: Ionic hydrides are crystalline solids and are non-volatile and non-conductive in the solid state. They conduct electricity only in the molten state or in solution, as they dissociate into ions.
32. The process of hydrogenation of vegetable oils requires a catalyst, typically: (c) Nickel
    • Explanation: Nickel is a commonly used catalyst for the hydrogenation of vegetable oils to produce vanaspati ghee.
33. Which property of hydrogen peroxide makes it useful as a bleaching agent? (a) Its strong oxidizing nature
    • Explanation: Hydrogen peroxide acts as a bleaching agent by releasing nascent oxygen (H2O2 → H2O + [O]), which oxidizes and thus bleaches colored substances.
34. The reaction of dihydrogen with oxygen to form water is: (b) Exothermic
    • Explanation: The formation of water from hydrogen and oxygen is a highly exothermic reaction, releasing a large amount of heat.

Part B: S-Block Elements

36. Which alkali metal has the highest hydration enthalpy? (a) Lithium
    • Explanation: Due to its very small ionic size and high charge density, Li+ ion is most intensely hydrated, hence has the highest hydration enthalpy among alkali metals.
37. Which of the following alkaline earth metals does not impart color to the flame? (d) Magnesium
    • Explanation: Beryllium and Magnesium do not show flame coloration. This is due to their small size and high ionization energies, requiring more energy to excite their electrons than the flame provides.
38. The general electronic configuration of alkali metals is: (a) ns1
    • Explanation: Alkali metals belong to Group 1, and their outermost shell contains one electron in the s-orbital.
39. Which of the following compounds is prepared by the Solvay process? (c) Na2CO3
    • Explanation: The Solvay process is an industrial method for the production of sodium carbonate (washing soda). Sodium bicarbonate (NaHCO3) is an intermediate product.
40. Anomalous behavior of Lithium is due to: (a) Its small size and high polarizing power
    • Explanation: Lithium’s anomalous properties (e.g., forming monoxide, reacting slowly with water, forming covalent compounds) are primarily attributed to its exceptionally small size and high charge density, leading to high polarizing power.
41. Which of the following is correct regarding the density of alkali metals down the group? (c) Increases, but Potassium is lighter than Sodium
    • Explanation: Generally, density increases down the group. However, potassium is an exception; its atomic volume is larger than expected, making it lighter than sodium.
42. Which of the following is an amphoteric oxide? (c) BeO
    • Explanation: Beryllium oxide (BeO) and Beryllium hydroxide (Be(OH)2) are amphoteric, meaning they can react with both acids and bases. Oxides of other alkaline earth metals are basic.
43. Which pair of elements shows a diagonal relationship? (c) Li and Mg
    • Explanation: Lithium shows a diagonal relationship with Magnesium, exhibiting similar properties due to their comparable charge-to-size ratios.
44. In the Castner-Kellner process, what is the product formed at the mercury cathode? (c) Sodium amalgam
    • Explanation: In the Castner-Kellner process for NaOH production, Na+ ions are discharged at the mercury cathode to form sodium amalgam (Na(Hg)).
45. The reaction of alkali metals with liquid ammonia produces a deep blue solution. This color is due to: (b) Ammoniated electrons
    • Explanation: When alkali metals dissolve in liquid ammonia, they produce solvated (ammoniated) metal ions and solvated (ammoniated) electrons, [e(NH3)y]-. It is these free, ammoniated electrons that absorb energy in the visible region and impart the deep blue color.
46. **When gypsum (CaSO4.2H2O) is heated to 393K, it forms: ** (c) Plaster of Paris
    • Explanation: Heating gypsum to 393K (120°C) results in the partial dehydration, forming calcium sulfate hemihydrate (CaSO4.1/2H2O), commonly known as Plaster of Paris.
47. Which alkali metal forms a peroxide when reacted with oxygen? (b) Sodium
    • Explanation: Lithium forms monoxide (Li2O). Sodium forms peroxide (Na2O2). Potassium, Rubidium, and Cesium form superoxides (MO2).
48. Which of the following is not a property of alkaline earth metals? (d) Generally form highly soluble hydroxides
    • Explanation: The solubility of alkaline earth metal hydroxides increases down the group, but they are generally less soluble than alkali metal hydroxides. Be(OH)2 and Mg(OH)2 are sparingly soluble or insoluble.
49. Which of the following is a constituent of Duralumin alloy? (b) Copper
    • Explanation: Duralumin is an aluminum alloy containing aluminum (main component), copper, magnesium, and manganese.
50. Quicklime is chemically: (c) Calcium oxide
    • Explanation: Quicklime is the common name for calcium oxide (CaO), produced by calcination of limestone.
51. The melting and boiling points of alkali metals decrease down the group. This is because: (b) Metallic bonding weakens
    • Explanation: As atomic size increases down the group, the strength of the metallic bonding (attraction between metal ions and delocalized electrons) decreases, leading to lower melting and boiling points.
52. Which of the following nitrates of alkaline earth metals decomposes on heating to give NO2 and O2? (d) All of the above
    • Explanation: Nitrates of alkaline earth metals (and most other metals, except Group 1 alkali metals) decompose on heating to give metal oxide, nitrogen dioxide (NO2), and oxygen (O2). The thermal stability of these nitrates increases down the group.
53. Alkali metals typically exhibit an oxidation state of: (a) +1
    • Explanation: Alkali metals have one valence electron (ns1) and readily lose it to form a stable +1 ion.
54. Which of the following is true about the basicity of hydroxides of alkaline earth metals down the group? (a) Increases
    • Explanation: As atomic size increases down Group 2, the ionic character of M-OH bond increases, and the strength of the M-OH bond decreases, making it easier for OH- to be released. Thus, basicity increases down the group (Be(OH)2 is amphoteric, Mg(OH)2 is a weak base, Ba(OH)2 is a strong base).
55. Which of the following is used in the manufacturing of cement? (d) All of the above
    • Explanation: Limestone (CaCO3) is a primary raw material, Quicklime (CaO) is an intermediate, and Gypsum (CaSO4.2H2O) is added to regulate the setting time of cement.
56. The solubility of alkaline earth metal sulfates decreases down the group. This is due to: (b) Decrease in hydration enthalpy
    • Explanation: For sulfates, the lattice enthalpy remains relatively constant down the group (due to large sulfate ion). However, the hydration enthalpy of the metal ions (M2+) decreases significantly down the group. Since hydration enthalpy decreases more rapidly than lattice enthalpy, the overall solubility decreases.
57. When hydrated lithium chloride (LiCl.2H2O) is heated, it forms: (b) Li2O
    • Explanation: Unlike other alkali metal halides, hydrated lithium chloride cannot be dehydrated to anhydrous LiCl by heating. Instead, it hydrolyzes and then decomposes to form lithium oxide (Li2O) and HCl. LiCl.2H2O -> LiOH + HCl + H2O. 2LiOH -> Li2O + H2O.
58. Which of the following alkaline earth metal salts is covalent? (c) BeCl2
    • Explanation: Due to its very small size and high polarizing power, Beryllium forms compounds with significant covalent character, such as BeCl2. Other alkaline earth metal chlorides are predominantly ionic.
59. The reducing character of alkali metals in aqueous solution follows the order: (a) Li > Na > K > Rb > Cs
    • Explanation: While Cs is the strongest reducing agent in the gaseous state, Lithium is the strongest reducing agent in aqueous solution. This is because of its exceptionally high hydration enthalpy, which more than compensates for its high ionization enthalpy.
60. Slaked lime is chemically: (c) Ca(OH)2
    • Explanation: Slaked lime is the common name for calcium hydroxide (Ca(OH)2), formed by adding water to quicklime (CaO).
61. Which of the following alkaline earth metal carbonates is most thermally stable? (d) BaCO3
    • Explanation: The thermal stability of alkaline earth metal carbonates increases down the group (BeCO3 < MgCO3 < CaCO3 < SrCO3 < BaCO3). This is because the increasing size of the metal ion stabilizes the larger carbonate ion, leading to higher lattice energies and greater stability.
62. The biological importance of sodium ions is related to: (d) All of the above
    • Explanation: Sodium ions play crucial roles in maintaining fluid balance, nerve impulse transmission, and muscle contraction in biological systems.
63. Which of the following properties is characteristic of both Lithium and Magnesium? (d) All of the above
    • Explanation: Due to diagonal relationship, Li and Mg both form stable nitrides (Li3N, Mg3N2), have covalent chlorides that are soluble in ethanol, and their hydroxides are weak bases.
64. Which of the following elements forms a superoxide on reaction with oxygen? (b) K
    • Explanation: Only Potassium (K), Rubidium (Rb), and Cesium (Cs) form superoxides (MO2) with oxygen. Lithium forms monoxide, and Sodium forms peroxide.
65. The alkaline earth metal that does not react with water even at high temperatures is: (c) Beryllium
    • Explanation: Beryllium is very unreactive with water, even at high temperatures, due to its small size and high ionization energy, and the formation of a protective oxide layer. Magnesium reacts slowly, while Calcium, Strontium, and Barium react increasingly vigorously.
66. Plaster of Paris is chemically: (b) Calcium sulfate hemihydrate
    • Explanation: Plaster of Paris is CaSO4.1/2H2O.
67. Which of the following statements about the solubility of alkali metal halides (LiF, NaCl, KBr, CsI) in water is generally true? (c) Follows no regular trend
    • Explanation: The solubility of alkali metal halides depends on a balance between lattice enthalpy and hydration enthalpy. There isn’t a simple monotonic trend down the group for all halides. For example, LiF has low solubility due to high lattice energy, while CsI also has low solubility due to low hydration energy.
68. The industrial preparation of sodium hydroxide is by the: (c) Castner-Kellner process
    • Explanation: The Castner-Kellner process is the electrolytic method used for the industrial production of sodium hydroxide (caustic soda) from brine.
69. Which of the following compounds of alkaline earth metals is used as an antacid? (b) Magnesium hydroxide
    • Explanation: Magnesium hydroxide (Mg(OH)2) is a common antacid, sold as Milk of Magnesia, due to its mild basicity.
70. The flame color produced by Barium is: (c) Apple green
    • Explanation: Calcium gives a brick-red flame, Strontium gives a crimson flame, and Barium gives an apple-green flame.