Position of Hydrogen in the Periodic Table
Hydrogen is the lightest and most abundant element in the universe. Its atomic number is 1 and electronic configuration is 1s¹. Placing hydrogen in the periodic table is tricky because it resembles both Group 1 (alkali metals) and Group 17 (halogens) in different ways.
Resemblance to Alkali Metals (Group 1)
- Both have one electron in the outermost s orbital.
- Both form monopositive ions (H⁺, Na⁺).
- Both have an oxidation state of +1 in most compounds.
- Both react with electronegative elements to form similar compounds (HCl like NaCl).
Resemblance to Halogens (Group 17)
- Both need one electron to complete their valence shell (H needs 1 to achieve He configuration).
- Both form diatomic molecules (H₂, Cl₂).
- Both can form a univalent anion (H⁻ like Cl⁻).
- Both form covalent compounds with non-metals (HCl, CCl₄).
Why Hydrogen is Unique
Despite these similarities, hydrogen is fundamentally different from both groups. It is a non-metal, it is a gas at room temperature, and it has properties that are unlike any other element. Hydrogen is therefore placed separately at the top of the periodic table.
Isotopes of Hydrogen
Hydrogen has three naturally occurring isotopes. Their properties differ more than the isotopes of any other element because the mass difference between them is proportionally very large (H has mass 1, D has mass 2 — a 100% difference, compared to ³⁵Cl vs ³⁷Cl which differ by only 6%).
| Isotope | Symbol | Protons | Neutrons | Abundance |
|---|---|---|---|---|
| Protium | ¹H (H) | 1 | 0 | 99.985% |
| Deuterium | ²H (D) | 1 | 1 | 0.015% |
| Tritium | ³H (T) | 1 | 2 | Trace (radioactive) |
Tritium is radioactive (β emitter) with a half-life of about 12.3 years. It is produced naturally by cosmic ray bombardment of nitrogen in the upper atmosphere and also in nuclear reactors.
Dihydrogen: Preparation and Properties
Laboratory Preparation
- Reaction of zinc with dilute sulphuric acid:
Zn + H₂SO₄ → ZnSO₄ + H₂↑ - Reaction of sodium hydride with water:
NaH + H₂O → NaOH + H₂↑ - Electrolysis of acidified water: H₂ is released at the cathode.
Industrial Preparation
- Steam reforming of natural gas:
CH₄ + H₂O (1270 K, catalyst) → CO + 3H₂ (water gas shift reaction gives more H₂) - Bosch process (water gas shift):
CO + H₂O → CO₂ + H₂ (gives high-purity H₂ for industrial use) - Electrolysis of brine:
H₂ released at the cathode, Cl₂ at the anode.
Physical Properties of Dihydrogen
H₂ is a colourless, odourless, tasteless gas. It is the lightest gas known. It has very low boiling point (−253 °C = 20 K) and melting point (−259 °C = 14 K). It is almost insoluble in water.
Chemical Properties
- Combustion: 2H₂ + O₂ → 2H₂O (highly exothermic; forms an explosive mixture with air). The oxyhydrogen flame (2000 °C) is used for welding.
- Reduction: H₂ reduces metal oxides to the metal. CuO + H₂ → Cu + H₂O.
- Reaction with halogens: H₂ + Cl₂ → 2HCl (fast with Cl₂; explosive with F₂; slow with Br₂; very slow with I₂).
- Reaction with alkali metals: 2Na + H₂ → 2NaH (sodium hydride); H₂ is the oxidising agent here (unusual for H₂).
Hydrides
Classify each hydride as ionic, covalent, or metallic. Reasoning is shown after your answer.
Ionic
Group 1 and 2 metals (s-block). H is H⁻ (hydride ion). React with water to give H₂. Conduct when molten.
Covalent
p-block elements. Shared electron pair bonds. Mostly gases or volatile liquids. Some are acidic (HCl), basic (NH₃), or neutral (CH₄).
Metallic / Interstitial
d-block transition metals. H fills interstices in metal lattice. Non-stoichiometric. Conduct electricity. Used for H₂ storage.
NaH
Sodium hydride
MgH₂
Magnesium hydride
CaH₂
Calcium hydride
AlH₃
Aluminium hydride
SiH₄
Silane
CH₄
Methane
NH₃
Ammonia
H₂O
Water
HF
Hydrogen fluoride
PdH₀.₆
Palladium hydride
TiH₂
Titanium hydride
LiAlH₄
Lithium aluminium hydride
B₂H₆
Diborane
GeH₄
Germane
SnH₄
Stannane
Compounds of hydrogen with other elements are called hydrides. There are three types:
| Type | Formation | Character | Examples |
|---|---|---|---|
| Ionic (saline) hydrides | Highly electropositive metals (Group 1, Ca, Sr, Ba) | Contain H⁻ ion; crystalline; react with H₂O to give H₂ | NaH, LiH, CaH₂ |
| Covalent (molecular) hydrides | Non-metals (p-block) | Covalent bond; vary from electron-deficient to electron-rich | CH₄, NH₃, H₂O, HF, B₂H₆ |
| Metallic (interstitial) hydrides | d-block and f-block metals | H occupies interstitial holes in metal lattice; non-stoichiometric; metallic luster | PdHₓ, TiH₂, LaH₂.₈₇ |
Ionic hydrides react with water vigorously: NaH + H₂O → NaOH + H₂↑. The hydride ion (H⁻) is a strong base.
Among covalent hydrides, those of Group 15-17 elements (NH₃, H₂O, HF) can form hydrogen bonds and are thus less volatile than expected. Boron hydride (B₂H₆, diborane) is electron-deficient and has unusual three-centre two-electron bonds.
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Water: Structure and Anomalous Properties
Hydrogen bonding occurs when H is bonded to a highly electronegative atom (F, O, N). It causes anomalously high boiling points.
Boiling points (Group 16 hydrides)
H₂Te
-2°C
H₂Se
-41°C
H₂S
-60°C
H₂O
+100°C
H-bond
Bar extends left for negative BP, right for positive. Zero line is in the middle.
Why H₂O and HF have anomalously high BPs
In Group 16, the expected trend from van der Waals forces would be BP increasing with molecular mass: H₂S < H₂Se < H₂Te. H₂O should have the lowest BP in the group — but it has +100°C, the highest! This is due to extensive hydrogen bonding: O has high electronegativity (3.44) and two lone pairs, so each H₂O forms up to 4 hydrogen bonds in liquid water, creating a strongly associated structure.
Same anomaly occurs with HF in Group 17 (F is the most electronegative element, 3.98). NH₃ also shows this anomaly in Group 15.
Hydrogen bonding conditions
1. H must be bonded to a highly electronegative atom: F, O, or N only.
2. The electronegative atom (donor) must have lone pairs to accept.
3. The bond is X-H...Y where X = F, O, N (donor); Y = F, O, N (acceptor).
Intermolecular H-bond: between two different molecules (water, HF, NH₃).
Intramolecular H-bond: within the same molecule (o-nitrophenol). Makes MP/BP lower than intermolecular.
Water (H₂O) is the most abundant compound on Earth. Its unique properties make it the solvent of life.
Structure of Water
Oxygen in water is sp³ hybridised. Two of the four sp³ orbitals form O-H bonds; the other two are occupied by lone pairs. The bond angle in water is 104.5° (less than the tetrahedral 109.5° because lone pair-lone pair repulsion is greater than bond pair-bond pair repulsion).
Water is a bent (V-shaped) polar molecule with a dipole moment of 1.84 D.
Hydrogen Bonding in Water
Each water molecule can form up to 4 hydrogen bonds: 2 as a hydrogen bond donor (O-H---O) and 2 as a hydrogen bond acceptor (via the lone pairs). This extensive 3D network of hydrogen bonds is responsible for all the anomalous properties of water.
Anomalous Properties of Water
- High boiling point (100 °C): Expected from molecular mass should be much lower. Compare H₂S (−60 °C). The extensive H-bonding requires a lot of energy to break.
- High heat of vaporisation (44 kJ/mol): Energy needed to break H-bonds.
- High surface tension: H-bonding creates strong intermolecular forces at the surface.
- Maximum density at 4 °C: Below 4 °C, water expands as it cools toward the more open ice structure. Above 4 °C, normal thermal expansion occurs. Maximum density at 4 °C = 1.0 g/cm³.
- Ice is less dense than liquid water: Ice has a more open, hexagonal lattice with more H-bonds per molecule than liquid water. Density of ice = 0.917 g/cm³ (lower than water at 4 °C). This is why ice floats on water.
- High dielectric constant (80): Makes water an excellent solvent for ionic compounds (it stabilises ions by solvation).
Chemical Properties of Water
- Amphoteric nature: Acts as an acid (H₂O ⇌ H⁺ + OH⁻) or as a base (H₂O + H⁺ → H₃O⁺).
- Hydrolysis: Reacts with many metal oxides and non-metal oxides to form acids or bases.
- Redox reactions: Water can be oxidised (to O₂) or reduced (to H₂) under extreme conditions.
Hydrogen Peroxide (H₂O₂)
Hydrogen peroxide is a pale blue (almost colourless) liquid. It is commercially available as a 30% aqueous solution (30 volumes = 30 times its volume of O₂ released on decomposition) and as 3% solution for antiseptic use.
Structure of H₂O₂
The O-O bond in H₂O₂ is a single bond (not a double bond as in O₂). The molecule has a non-planar (open-book) structure. The O-O-H bond angle is about 96.7°. The dihedral angle between the two O-H bonds is about 111.5° in the gas phase. The O-O bond length is 1.47 Å.
Preparation
- Industrial: electrolysis of 50% H₂SO₄ to give peroxodisulphuric acid, then hydrolysis.
- Modern industrial: auto-oxidation of 2-ethylanthrahydroquinone in the presence of O₂.
- Lab: BaO₂ + H₂SO₄ → BaSO₄↓ + H₂O₂ (barium peroxide with dilute H₂SO₄).
- Na₂O₂ + H₂SO₄ → Na₂SO₄ + H₂O₂.
Properties of H₂O₂
H₂O₂ decomposes slowly: 2H₂O₂ → 2H₂O + O₂. Decomposition is accelerated by light, heat, and catalysts (MnO₂, Fe²⁺). It is stored in dark polyethylene bottles with a small amount of urea or phosphoric acid to slow decomposition.
Oxidising Agent (more common role)
Reducing Agent (in the presence of stronger oxidising agents)
Uses of H₂O₂
- Bleaching agent for textiles, paper, hair.
- Antiseptic (3% solution).
- As a rocket propellant (concentrated).
- Manufacturing of epoxy resins and pharmaceuticals.
- Treatment of sewage and polluted water.
Heavy Water (D₂O)
Heavy water is water in which the hydrogen atoms are replaced by deuterium (²H or D). Its formula is D₂O. It is slightly different from ordinary water in its physical properties.
| Property | H₂O | D₂O |
|---|---|---|
| Boiling point | 100 °C | 101.4 °C |
| Melting point | 0 °C | 3.8 °C |
| Density (at 25 °C) | 0.997 g/cm³ | 1.105 g/cm³ |
| Temperature of max density | 4 °C | 11.6 °C |
D₂O is used as a moderator in nuclear reactors (to slow neutrons) and as a tracer in studying reaction mechanisms.
D₂O is toxic in large amounts (it slows biological reactions; organisms eventually die if D₂O completely replaces H₂O in their cells).
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Worked NEET Problems
NEET-style problem · Isotopes of Hydrogen
Question
Which isotope of hydrogen is radioactive? What is its symbol and approximate half-life?
Solution
Tritium (³H or T) is the radioactive isotope of hydrogen. It decays by beta emission (emits an electron from the nucleus, so a neutron converts to a proton). Its half-life is approximately 12.3 years.
Protium (¹H) and deuterium (²H) are stable (non-radioactive).
NEET-style problem · H₂O₂ as Oxidising and Reducing Agent
Question
In which reaction does H₂O₂ act as a reducing agent? (a) H₂O₂ + 2KI → I₂ + 2KOH (b) 2KMnO₄ + 5H₂O₂ + 3H₂SO₄ → 2MnSO₄ + K₂SO₄ + 5O₂ + 8H₂O
Solution
In reaction (a), H₂O₂ oxidises I⁻ to I₂ — here H₂O₂ is the oxidising agent.
In reaction (b), H₂O₂ reduces KMnO₄ (Mn goes from +7 to +2). H₂O₂ itself is oxidised (O in H₂O₂ goes from −1 to 0 in O₂). So H₂O₂ acts as the reducing agent in reaction (b).
NEET-style problem · Anomalous Properties of Water
Question
Why is water called the universal solvent?
Solution
Water dissolves more substances than any other common solvent because of two main features:
(1) High dielectric constant (about 80): It greatly reduces the electrostatic attraction between ions, allowing ionic compounds to dissociate easily.
(2) Ability to form hydrogen bonds: Water can hydrogen-bond with polar molecules (like sugars and proteins), solvating them and bringing them into solution.
Its ability to dissolve ionic, polar covalent, and even some non-polar substances (when aided by surfactants) earns it the title of universal solvent.
Summary Cheat Sheet
| Concept | Key Point |
|---|---|
| Hydrogen's position | Resembles Group 1 (loses e⁻) and Group 17 (gains e⁻); placed separately at top |
| Protium / Deuterium / Tritium | ¹H / ²H (D) / ³H (T); T is radioactive (β emitter, t½ ≈ 12.3 yr) |
| Ionic hydrides | H⁻ formed with Group 1 & 2 metals; react with water to give H₂ |
| Metallic hydrides | d-block/f-block metals; non-stoichiometric; H in interstitial holes |
| Water bond angle | 104.5° (sp³ O with 2 lone pairs) |
| Ice floats on water | Open hexagonal H-bonded lattice → ice less dense than liquid water |
| Max density of water | 4 °C = 1.0 g/cm³ |
| H₂O₂ structure | Non-planar (open-book); O-O single bond (1.47 Å); dihedral ≈ 111.5° |
| H₂O₂ oxidising agent | Oxidises Fe²⁺, I⁻, PbS; O: −1 → −2 |
| H₂O₂ reducing agent | Reduces KMnO₄, Cl₂; O: −1 → 0 (in O₂) |
| H₂O₂ storage | Dark bottles; stabilised with urea/H₃PO₄; light/MnO₂ accelerate decomposition |
| D₂O | Bp 101.4 °C; mp 3.8 °C; density 1.105 g/cm³; used as nuclear moderator |
Frequently asked questions
Why does hydrogen have a unique position in the periodic table?
Hydrogen resembles Group 1 (alkali metals) in that it has one electron in its outermost shell and tends to form H⁺ by losing that electron. It resembles Group 17 (halogens) in that it needs one more electron to complete its valence shell and can form H⁻ (hydride ion). However, hydrogen differs fundamentally from both groups: it is a non-metal, forms a diatomic molecule (H₂), and has a very low ionisation energy compared to halogens but does not form salts the way alkali metals do. So it occupies a unique, anomalous position.
Why does ice float on water?
In ice, water molecules form a hydrogen-bonded lattice with an open, cage-like hexagonal structure. This structure has larger intermolecular spaces than liquid water, making ice less dense than liquid water (density of ice ≈ 0.917 g/cm³ vs water = 1.0 g/cm³). When ice melts, some hydrogen bonds break and the open structure collapses, allowing molecules to pack more closely. This is why water has maximum density at 4 °C, not at 0 °C.
What are the oxidising and reducing properties of H₂O₂?
H₂O₂ can act as both an oxidising agent and a reducing agent. As an oxidising agent (more common): H₂O₂ + 2H⁺ + 2e⁻ → 2H₂O (gains electrons, oxidises the substrate). Example: oxidises Fe²⁺ to Fe³⁺, and bleaches coloured matter. As a reducing agent: H₂O₂ → O₂ + 2H⁺ + 2e⁻ (loses electrons when reacting with a stronger oxidising agent). Example: reduces Cl₂ to HCl, and reduces acidic KMnO₄.
What are the three types of hydrides?
(1) Ionic (saline) hydrides: formed by highly electropositive metals (Group 1 and some Group 2). The H is H⁻ (hydride ion). Examples: NaH, CaH₂. They are crystalline solids that react vigorously with water to give H₂. (2) Covalent (molecular) hydrides: formed by non-metals. H is covalently bonded. Examples: CH₄, NH₃, H₂O, HF. (3) Metallic (interstitial) hydrides: formed by transition metals (d-block) and lanthanides/actinides. H atoms occupy interstitial positions in the metal lattice. Non-stoichiometric. Examples: PdHₓ, TiH₂.
Why does water have an anomalously high boiling point for its molecular mass?
Water (M = 18 g/mol) has a boiling point of 100 °C, which is much higher than expected by comparison with other Group 16 hydrides (H₂S: −60 °C, H₂Se: −41 °C). The reason is extensive hydrogen bonding between water molecules. Each water molecule can form up to 4 hydrogen bonds (2 as donor, 2 as acceptor). A lot of energy is needed to break these hydrogen bonds before vaporisation can occur, resulting in a high boiling point.
What is the structure of H₂O₂ and what does it mean that it is non-planar?
H₂O₂ has an open-book (butterfly) structure. The two O-H bonds are not in the same plane; the dihedral (torsional) angle between the two O-H bonds is about 111.5° in the gas phase (different in solid due to hydrogen bonding). The O-O bond is a single bond (longer than the O=O in O₂). The non-planar structure means H₂O₂ is chiral in principle, but the two mirror-image forms interconvert rapidly.
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