%20and%20Properties.webp "Classification of Hydrides: Ionic, Covalent and Metallic Hydrides in Inorganic Chemistry")
Chapter 8: Hydrogen and the Hydrides — A Complete Study Guide for JEE, NEET, IIT-JAM & BITSAT
Hydrogen is one of the most fascinating elements in the entire periodic table. It holds the distinction of being the simplest, lightest, and most abundant element in the universe. Despite its simplicity, hydrogen's chemistry is surprisingly rich — it forms three different types of ions, three different classes of hydrides, and plays a central role in acid-base chemistry. This chapter is essential for competitive exams because it cuts across multiple concepts: electronic structure, bonding, industrial processes, isotopes, and theories of acids and bases.
Let's break this chapter down clearly, step by step, so that even if you're encountering it for the first time, you walk away with a solid and exam-ready understanding.
1. Electronic Structure of Hydrogen JEE | NEET
Hydrogen has the electronic configuration 1s1 — a nucleus with a single proton (+1 charge) surrounded by one electron in the 1s orbital. This means hydrogen can achieve stability in three different ways:
- By forming a covalent bond — sharing its electron with another atom (e.g., H2, H2O, HCl, CH4). This is the most common method.
- By losing its electron to form H+ — the proton (H+) is extremely small (radius ≈ 1.5 × 10−5 Å, compared to 0.74 Å for the H atom) and has very high polarizing power. Free protons do NOT exist in normal conditions — they are always associated with other molecules. In water, H+ exists as H3O+ or H9O4+.
- By gaining an electron to form H− — the hydride ion, found in ionic hydrides like LiH. This is uncommon and only happens with highly electropositive metals (Group 1, some Group 2).
2. Position of Hydrogen in the Periodic Table NEET | BITSAT
Hydrogen's position in the periodic table is genuinely unique and debated. It shares characteristics with three different groups:
| Group Resemblance | Reason | Key Difference |
|---|---|---|
| Group 1 (Alkali Metals) | Both have 1 electron in outer shell | Alkali metals lose e⁻ easily; H prefers covalent bonds |
| Group 17 (Halogens) | Both are one electron short of noble gas structure | Halogens gain e⁻ readily; H does not typically form H⁻ |
| Group 14 (Carbon group) | Both have a half-filled outer shell | H has only one electron; C has four valence electrons |
3. Abundance of Hydrogen
Hydrogen is the most abundant element in the universe — approximately 92% of all atoms in the universe are hydrogen. Yet, in the Earth's atmosphere, free H2 is very rare because Earth's gravity is too weak to hold such a light gas. However, hydrogen is the tenth most abundant element in Earth's crust (1520 ppm or 0.152% by weight). It is found in vast quantities in:
- Water (oceans, rivers, ice caps)
- Living matter — carbohydrates, proteins
- Fossil fuels — coal, petroleum, natural gas
- Ammonia, acids, and thousands of organic compounds
4. Preparation of Hydrogen JEE | IIT-JAM
Hydrogen is manufactured industrially and prepared in the lab by a variety of methods. Let's understand each one with the exact chemistry:
Method 1: Steam over Red Hot Coke (Water Gas)
Water gas (CO + H₂) is an important industrial fuel. To separate H2, CO can be removed by liquefying it or by passing over iron oxide with steam at 450°C (shift converter):
Method 2: Steam Reformer Process
Methane (natural gas) is mixed with steam and passed over a nickel catalyst at 800–900°C:
This hydrogen is used in the Haber process (making NH3) and for hardening oils.
Method 3: Oil Refinery Cracking
High molecular weight hydrocarbons (naphtha, fuel oil) are cracked to produce petrol. Hydrogen is a valuable by-product.
Method 4: Electrolysis of Water (Purest Method)
Very pure hydrogen (99.9%) is obtained by electrolysis of NaOH or KOH solution:
This is the most expensive method.
Method 5: Chlor-Alkali Industry
Electrolysis of aqueous NaCl produces NaOH, Cl2, and H2 as a by-product.
Method 6: Laboratory Method
Method 7: Reaction of Salt-like Hydrides with Water
5. Properties of Molecular Hydrogen JEE | NEET
Hydrogen gas (H2) is colourless, odourless, and almost insoluble in water. It is the lightest gas known, which is why it is used (instead of helium) to fill meteorological balloons.
- H2 forms diatomic molecules with a very strong H–H covalent bond (bond energy = 435.9 kJ mol−1).
- Hydrogen is not very reactive under normal conditions — this is due to kinetics (the high activation energy needed to break the H–H bond), not thermodynamics.
- Many reactions of H2 require high temperatures, high pressures, or catalysts (often transition metals).
Key Reactions of Molecular Hydrogen
1. Combustion in Air/Oxygen:
This reaction releases a large amount of energy. Mixtures of H2 and O2 near 2:1 ratio are explosive. Temperature of ~3000°C can be achieved (used in oxy-hydrogen welding and cutting metals).
2. Reaction with Halogens:
3. Haber Process (making ammonia):
Conditions: Fe catalyst, 380–450°C, 200 atmospheres pressure.
4. Hydrogenation of Oils (making margarine):
Palladium catalyst; converts unsaturated fatty acids (liquid oils) to saturated fats (solid at room temperature).
5. Production of Methanol:
6. Isotopes of Hydrogen JEE | NEET | IIT-JAM
Hydrogen has three naturally occurring isotopes — one of the most important distinguishing features of this element in chemistry exams.
| Isotope | Symbol | Protons | Neutrons | Abundance | Radioactive? |
|---|---|---|---|---|---|
| Protium | ¹H (H) | 1 | 0 | 99.986% | No |
| Deuterium | ²H (D) | 1 | 1 | 0.014% | No |
| Tritium | ³H (T) | 1 | 2 | 7 × 10⁻¹⁶ % | Yes (β emission) |
Key Physical Data (Table 8.1 equivalent)
| Property | H₂ | D₂ | T₂ |
|---|---|---|---|
| Boiling Point (°C) | −252.6 | −249.3 | −248.0 |
| Bond Length (Å) | 0.7414 | 0.7414 | (0.7414) |
| Heat of Dissociation (kJ mol⁻¹) | 435.9 | 443.4 | 446.9 |
Isotope Effects
Because hydrogen isotopes differ greatly in mass (H=1, D=2, T=3), they show larger differences in physical and chemical properties than isotopes of any other element. This is called the isotope effect.
- H2 is adsorbed on surfaces more rapidly than D2
- H2 reacts 13 times faster with Cl2 than D2 (lower activation energy)
- Protium water (H2O) dissociates about 3 times more than heavy water (D2O)
Tritium — The Radioactive Isotope
Tritium is produced in nuclear reactors by irradiating lithium with slow neutrons:
Uses: thermonuclear devices, radioactive tracer in research (emits only low-energy β radiation, no shielding required).
Heavy Water (D₂O)
| Property | H₂O | D₂O |
|---|---|---|
| Freezing Point (°C) | 0 | 3.82 |
| Boiling Point (°C) | 100 | 101.42 |
| Density at 20°C (g cm⁻³) | 0.917 | 1.017 |
| Ionic product Kw (25°C) | 1.0 × 10⁻¹⁴ | 3.0 × 10⁻¹⁵ |
D2O is prepared by electrolysis of water — protium bonds break more readily, so H2 is liberated faster, and D2O remains. About 29,000 litres of water must be electrolysed to get 1 litre of 99% pure D2O.
7. Ortho and Para Hydrogen IIT-JAM | JEE Advanced
The H2 molecule exists in two forms based on the relative spin of the two proton nuclei:
- Ortho-H₂: Both nuclear spins are parallel (same direction) — higher energy form.
- Para-H₂: Nuclear spins are antiparallel (opposite directions) — lower energy form.
- At absolute zero: 100% para-H₂
- At high temperatures: ~75% ortho-H₂, 25% para-H₂
- Para-H₂ is prepared by passing the gas through charcoal cooled to liquid air temperature.
- Catalysts for ortho-para conversion: activated charcoal, Fe, Ni, Pt, W, O2, NO, NO2, Co2+, Cr2O3
8. Hydrides — Classification and Properties JEE | NEET | IIT-JAM
Binary compounds of elements with hydrogen are called hydrides. Based on the type of bonding (which depends on electronegativity), hydrides are classified into three main types:
8.1 Ionic (Salt-like) Hydrides
Formed by Group 1 alkali metals and heavier Group 2 metals (Ca, Sr, Ba) at high temperatures. Evidence for ionic nature:
- Molten LiH (m.p. 691°C) conducts electricity; H2 is liberated at the anode, confirming H− ion.
- CaH2 dissolves in eutectic LiCl/KCl melt; electrolysis evolves H2 at anode.
- Crystal structures show no directional bonding (ionic).
Important reactions of ionic hydrides:
8.2 Covalent Hydrides
Formed by p-block elements. They consist of discrete molecules held together by weak van der Waals forces, making them volatile with low melting and boiling points.
| Compound | m.p. (°C) | b.p. (°C) | Special Feature |
|---|---|---|---|
| CH₄ | −183 | −162 | No hydrogen bonding |
| NH₃ | −78 | −33 | H-bonding (b.p. unexpectedly high) |
| H₂O | 0 | +100 | Strong H-bonding (very high b.p.) |
| HF | −83 | +20 | H-bonding (zig-zag chains) |
| SiH₄ | −185 | −111 | No H-bonding |
| H₂S | −86 | −60 | Weak H-bonding |
The Group 13 hydrides are unusual — they are electron-deficient and polymeric. Diborane (B2H6) is the simplest boron hydride, with multi-centre bonding (discussed separately).
8.3 Metallic (Interstitial) Hydrides
Formed by d-block and f-block elements. They are non-stoichiometric (variable composition), e.g.:
- LaH2.87, TiH1.8, ZrH1.9, PdH0.7
The "hydrogen gap" refers to elements in the middle of the d-block (Mn, Fe, Co, Ni, etc.) that do NOT form stable hydrides. Metallic hydrides have properties similar to parent metals — metallic lustre, electrical conductivity, magnetic properties. They are less dense than the parent metal (crystal lattice expands).
The extraordinary Pd/H₂ system: Red hot palladium can absorb up to 935 times its own volume of H2. This is used to separate H2 from other gases.
9. Hydrogen Bonding JEE | NEET | BITSAT
A hydrogen bond is a weak electrostatic attraction between a lone pair of electrons on an electronegative atom (F, O, N, Cl) and a covalently bonded hydrogen atom carrying a partial positive charge (δ+).
Key Facts about Hydrogen Bonds
- Bond energy: typically 10–45 kJ mol−1 (compare: C–C covalent bond = 347 kJ mol−1)
- Despite being weak, they are critically important in biochemistry and materials science.
- The four most important electronegative atoms for H-bonding: F, O, N, Cl
Why Does Hydrogen Bond Form Only with F, O, N, Cl?
These elements are the most electronegative, and when bonded to hydrogen they create a large partial positive charge on H (δ+). This δ+ H is then attracted to lone pairs on adjacent electronegative atoms.
Effect on Physical Properties
From the graph, it's clear that H2O, NH3, and HF have anomalously high boiling points compared to other hydrides in their groups. Without hydrogen bonding, water would boil at about −100°C instead of +100°C — life as we know it would not exist.
Types of Hydrogen Bonding
- Intermolecular H-bonding: Between different molecules (e.g., in H₂O, HF, NH₃). Raises boiling point, melting point, and enthalpy of vaporization.
- Intramolecular H-bonding: Within the same molecule (e.g., in ortho-nitrophenol). Reduces boiling point compared to meta and para isomers and reduces acidity.
Biological Importance of Hydrogen Bonds
- Link polypeptide chains in proteins (secondary structure)
- Hold complementary base pairs together in DNA double helix
- Responsible for water being liquid at room temperature
- Have low activation energy → play important roles in reactions at normal temperatures
10. The Hydrogen Ion (H⁺) JEE | NEET
The ionization energy of hydrogen is 1311 kJ mol−1 — extremely large. Therefore, H+ ions (bare protons) do not exist freely. In water, H+ is hydrated to form:
- H3O+ (hydronium ion) — primary solvated form
- H9O4+ — extended hydration shell
The hydration energy of H+ in water is 1091 kJ mol−1, which largely offsets the ionization energy. The remaining ~220 kJ comes from the electron affinity of the anion and its solvation energy.
11. Acids and Bases — Comprehensive Theories JEE | NEET | IIT-JAM | BITSAT
This section is one of the most concept-heavy and frequently tested portions of this chapter. There are multiple theories, and each has its own scope and utility.
11.1 Arrhenius Theory
Proposed in 1884:
- Acid: Substance that produces H+ in water
- Base: Substance that produces OH− in water
Limitation: Only applicable to aqueous solutions.
The pH Scale
At 25°C: pure water has [H+] = [OH−] = 10−7 mol/L → pH = 7 (neutral)
Acid Strength Trends from pKₐ values:
| Compound | CH₄ | NH₃ | H₂O | HF | HCl | HBr | HI |
|---|---|---|---|---|---|---|---|
| pKₐ | 46 | 35 | 16 | 3 | −7 | −9 | −10 |
Acid strength increases across a period (left to right) and increases down a group (for binary hydrides). Lower pKₐ = stronger acid.
11.2 Brønsted–Lowry Theory (1923)
- Acid: Proton (H+) donor
- Base: Proton (H+) acceptor
Key concept — conjugate acid-base pairs:
A strong acid has a weak conjugate base, and vice versa.
Advantage of Brønsted–Lowry: Extends to non-aqueous solvents like liquid NH3, glacial acetic acid, anhydrous H2SO4.
Levelling and Differentiating Solvents:
- Levelling solvent (e.g., liquid NH3): Makes all strong acids appear equally strong by accepting their protons completely.
- Differentiating solvent (e.g., glacial acetic acid): Reveals differences in acid strength since even strong acids only partially ionize here.
11.3 Lewis Theory
- Lewis Acid: Electron pair acceptor (e.g., BF3, AlCl3, metal ions, CO2)
- Lewis Base: Electron pair donor (e.g., NH3, H2O, F−, CN−)
Limitation: No universal scale of acid/base strength; almost all reactions become acid-base reactions.
11.4 The Solvent System (Cady–Elsey)
The most general definition — applies to any self-ionizing solvent (regardless of whether it contains protons):
| Solvent | Self-Ionization | Acid species | Base species |
|---|---|---|---|
| H₂O | 2H₂O ⇌ H₃O⁺ + OH⁻ | H₃O⁺ | OH⁻ |
| NH₃ (liq.) | 2NH₃ ⇌ NH₄⁺ + NH₂⁻ | NH₄⁺ | NH₂⁻ |
| H₂SO₄ | 2H₂SO₄ ⇌ H₃SO₄⁺ + HSO₄⁻ | H₃SO₄⁺ | HSO₄⁻ |
| N₂O₄ | N₂O₄ ⇌ NO⁺ + NO₃⁻ | NO⁺ | NO₃⁻ |
11.5 Lux–Flood Theory
Used in high-temperature anhydrous systems (metallurgy, ceramics):
- Acid: Oxide that accepts O2− (e.g., CO2, SiO2)
- Base: Oxide that donates O2− (e.g., CaO, Na2O)
11.6 Hard and Soft Acids and Bases (HSAB Theory — Pearson)
Metal ions (Lewis acids) and ligands (Lewis bases) are classified as hard or soft based on polarizability:
| Hard (small, low polarizability) | Soft (large, high polarizability) | |
|---|---|---|
| Acids | H⁺, Li⁺, Na⁺, Mg²⁺, Al³⁺, BF₃, Cr³⁺ | Pd²⁺, Pt²⁺, Cu⁺, Ag⁺, Au⁺, Hg²⁺ |
| Bases | F⁻, OH⁻, NH₃, H₂O, Cl⁻, CO₃²⁻ | H⁻, CN⁻, SCN⁻, I⁻, CO, R₂S |
Pearson's Rule: Hard acids prefer to react with hard bases; soft acids prefer to react with soft bases.
12. The Hydrogen Economy
The concept of the hydrogen economy proposes replacing fossil fuels with hydrogen as the primary energy carrier. Key advantages:
- Burning H2 produces only water — no CO2, no SO2, no carcinogenic hydrocarbons
- Hydrogen can be produced by electrolysis and chemical methods
- Can be stored as gas, liquid (cryogenic flasks), or absorbed in metal alloys (e.g., LaNi5 absorbs 7 mol H2/mol alloy)
- Liquid H2 is used as rocket fuel (Saturn series, Space Shuttle)
However, there are challenges: explosive risk, cost of production, and storage difficulties.
13. Quick Revision: Most Important Points for Exams
- Electronic configuration of H: 1s¹. It can form H⁺, H⁻, or covalent bonds.
- Hydrogen is the most abundant element in the universe (92% of atoms).
- Three isotopes: Protium (¹H, 99.986%), Deuterium (²H, 0.014%), Tritium (³H — radioactive, β emitter, t½ = 12.26 years).
- H–H bond energy = 435.9 kJ mol⁻¹ (very strong; reason for low reactivity).
- D₂O has higher boiling point (101.42°C), melting point (3.82°C), and density (1.017 g/cm³) than H₂O.
- 29,000 L of water must be electrolysed for 1 L of 99% D₂O.
- Ortho-H₂ = parallel spins; Para-H₂ = antiparallel spins. Para is more stable (lower energy).
- Ionic hydrides react with water to give H₂. They conduct electricity in molten state.
- Metallic hydrides are non-stoichiometric. The hydrogen gap = elements in mid-d-block that don't form stable hydrides.
- Pd can absorb up to 935 times its own volume of H₂.
- H-bonding occurs with F, O, N, Cl. Bond energy: 10–45 kJ mol⁻¹.
- Without H-bonding, water would boil at −100°C.
- Brønsted–Lowry: acid = proton donor, base = proton acceptor.
- Lewis: acid = electron pair acceptor, base = electron pair donor.
- HSAB: Hard acids prefer hard bases; soft acids prefer soft bases.
- Lux–Flood: acid = oxide accepting O²⁻; base = oxide donating O²⁻ (used in metallurgy).
- Acid strength across period: CH₄ < NH₃ < H₂O < HF; down a group: HF < HCl < HBr < HI.
- NaBH₄ and LiAlH₄ are made from NaH — important reducing agents.
14. Exam-Style Practice Concepts
Q1 (NEET type): Which of the following is NOT correct about D₂O?
Answer key: D₂O has a lower dielectric constant than H₂O (80.5 vs 82 at 20°C), so ionic compounds are generally LESS soluble in D₂O than in H₂O.
Q2 (JEE type): Why is the boiling point of HF (+20°C) much higher than HCl (−85°C), even though F is in Period 2 and Cl in Period 3?
Answer key: HF molecules form strong intermolecular hydrogen bonds (H···F, bond energy ~29 kJ/mol) due to the high electronegativity of F and the small size of the F atom. HCl cannot form significant H-bonds.
Q3 (IIT-JAM type): Explain why Ag⁺ preferentially reacts with I⁻ and not F⁻.
Answer key: By HSAB theory, Ag⁺ is a soft acid and I⁻ is a soft base — they prefer each other. F⁻ is a hard base, which pairs better with hard acids like H⁺ or Na⁺.
Q4 (BITSAT type): The reaction of H₂ with Cl₂ is explosive in sunlight but slow in the dark. Why?
Answer key: Sunlight provides the energy to initiate a free radical chain reaction (photocatalysis). Cl₂ → 2Cl• (by UV), then Cl• + H₂ → HCl + H•, H• + Cl₂ → HCl + Cl•, etc. In the dark, there is insufficient energy to break Cl–Cl bond.
15. Chapter Summary at a Glance
Hydrogen is not just a simple element — it is the gateway to understanding bonding, industrial chemistry, acid-base theories, and even the future of clean energy. Every concept in this chapter connects deeply to real-world chemistry and competitive examination problems. Master this chapter well, and you'll find it rewards you with marks across multiple topics in JEE, NEET, IIT-JAM, and BITSAT.
%20and%20Properties.webp&description=Hydrogen and Hydrides: Complete Guide to Types (Ionic, Covalent, Metallic) and Properties In Inorganic Chemistry For IIT-JAM, BITSAT, NEET, JEE, TGT, PGT Exams){kind=link}