Course Contents
Introduction to Chemical Bonding
Chemical Bond: The force which holds two or more atoms or ions to form a large variety of compounds.
- Octet Rule: Every atom tries to attain electronic configuration near to a noble gas (maximum of eight electrons in valence shell).
- Energy Stability: Each atom tries to attain lowest energy state because it is a stable state.
Types of Bonds:
- Electrovalent or Ionic Bond
- Covalent Bond
- Co-ordinate Covalent or Dative Bond
Atoms bond to achieve noble gas configuration (octet) and lower their energy – both conditions must be satisfied!
Atomic Sizes
| Particulars | Atomic Radii | Ionic Radii | Covalent Radii |
|---|---|---|---|
| Definition | Average distance between nuclei of an atom and its outermost electronic shell | Average distance between nuclei of an ion and its outermost electronic shell | Half of the single bond length between two similar atoms covalently bonded |
| Trend along period | Decreases (left to right) | Decreases (left to right) | Decreases (left to right) |
| Trend down group | Increases (top to bottom) | Increases (top to bottom) | Increases (top to bottom) |
| Effect of shielding | Greater shielding = larger radii | Greater shielding = larger radii | Greater shielding = larger radii |
Size trends: Across period → decreases (more protons pull electrons), Down group → increases (more shells). Cations smaller, anions larger than parent atoms!
Ionization Energy, Electron Affinity & Electronegativity
| Particular | I.P/I.E (Ionization Potential/Energy) | E.A (Electron Affinity) | E.N (Electronegativity) |
|---|---|---|---|
| Definition | Energy required to remove an electron from outermost shell of gaseous atom | Energy released when electron is added to outermost shell of gaseous atom | Tendency of an atom to attract shared pair of electrons |
| First, second energies | 1st I.P is lesser, 2nd is higher, 3rd is much higher | 1st E.A is exothermic, 2nd and 3rd E.A are endothermic | Not applicable |
| Trend down group | Decreases (top to bottom) | Decreases (top to bottom) | Decreases (top to bottom) |
| Trend along period | Increases (left to right) | Increases (left to right) | Increases (left to right) |
| Effect of size | I.P decreases as size increases | E.A decreases as size increases | E.N decreases as size increases |
I.E = Energy to remove electron (harder for smaller atoms), E.A = Energy released adding electron (easier for smaller atoms), E.N = Ability to attract electrons!
Electrovalent or Ionic Bond
Definition: Bond formed by complete transfer of electrons from more electropositive to more electronegative elements.
- Electropositive elements (low I.E) like I-A and II-A that lose electrons to form cations
- Electronegative elements (high E.A) like VI-A and VII-A that gain electrons to form anions
Characteristic of Ionic Compounds:
- Good conductors in fused state/aqueous solution, poor in solid state
- Soluble in polar solvents, insoluble in non-polar solvents
- High melting and boiling points
- Form ionic crystals with orderly arrangement
- Hard and brittle
- Maximum ionic character is 92% observed in CsF
- ∆E.N of bonded atoms ∝ % ionic character
Cl(g) + e⁻ → Cl⁻(g) (E.A = -349 kJ/mole)
Na⁺(g) + Cl⁻(g) → NaCl(s)
Ionic bond = Metal + Non-metal = Complete electron transfer = High melting point + Conducts when melted/dissolved!
Covalent Bond
Definition: Bond formed by mutual sharing of electrons between two atoms.
Types of Covalent Bonds:
- Single: Sharing one electron from each atom (represented by —)
- Double: Sharing two electrons from each atom (represented by =)
- Triple: Sharing three electrons from each atom (represented by ≡)
Non-polar Covalent Bond
- Between similar atoms or negligible E.N difference (0-0.4)
- Electron cloud symmetrical
- No partial charges
- Weak intermolecular attraction
- Low boiling/melting points
- Example: Cl—Cl
Polar Covalent Bond
- Between dissimilar atoms with E.N difference (0.4-1.7)
- Electron cloud unsymmetrical
- Partial charges develop
- Strong intermolecular attraction
- High boiling/melting points
- Example: H—Cl
Covalent bond = Sharing electrons = Non-metal + Non-metal = Low melting point + Poor conductors (except graphite)!
Co-ordinate Covalent Bond (Dative Bond)
Definition: Covalent bond formed by donation of electron-pair by one bonded atom to another.
- Formed by overlapping of completely filled orbital with empty orbital
- Donor atom = Lewis base (has lone pair)
- Acceptor atom = Lewis acid (has empty orbital)
- Represented by arrow (→) pointing from donor to acceptor
- Single atom contributes to stability of both atoms
Examples:
- Ammonium ion (NH₄⁺): Nitrogen in NH₃ donates lone pair to H⁺
- NH₃ + BF₃: Nitrogen donates lone pair to electron-deficient boron
NH₃ + BF₃ → H₃N→BF₃
Dative bond = One-sided sharing = Lone pair donor + Empty orbital acceptor = Arrow from donor to acceptor!
VSEPR Theory
Postulates:
- Valence electron pairs (lone and bond pairs) arrange to maximize distance and minimize repulsion
- Lone pairs occupy more space than bond pairs
- Repulsion order: Lone-Lone > Lone-Bond > Bond-Bond
- Multiple bonds behave like single bonds in determining geometry
| Type | Total e⁻ pairs | Bonding | Lone | Arrangement | Molecular Geometry | Examples |
|---|---|---|---|---|---|---|
| AB₂ | 2 | 2 | 0 | Linear | Linear | BeCl₂, CO₂ |
| AB₃ | 3 | 3 | 0 | Trigonal planar | Trigonal planar | BF₃, SO₃ |
| AB₃ | 3 | 2 | 1 | Trigonal planar | Bent/Angular | SO₂, SnCl₂ |
| AB₄ | 4 | 4 | 0 | Tetrahedral | Tetrahedral | CH₄, CCl₄ |
| AB₄ | 4 | 3 | 1 | Tetrahedral | Trigonal pyramidal | NH₃, NF₃ |
| AB₄ | 4 | 2 | 2 | Tetrahedral | Bent/Angular | H₂O, H₂S |
VSEPR: Electron pairs repel → Lone pairs repel most → More lone pairs = More bent shape. Remember: CH₄ (0 lp) = Tetrahedral, NH₃ (1 lp) = Pyramidal, H₂O (2 lp) = Bent!
Sigma and Pi Bonds
Sigma (σ) Bond
- Formed by head-to-head overlapping
- Electron density concentrated around bond axis
- Only one lobe between nuclei
- Only one σ bond between two atoms
- Free rotation about bond axis
- No nodal plane
- Determines molecular shape
Pi (π) Bond
- Formed by parallel/sideways overlapping
- Electron density above and below bond axis
- Two lobes on opposite sides
- One or more π bonds between two atoms
- No free rotation about bond axis
- Has nodal plane along bond axis
- No effect on molecular shape
Examples:
- H₂: 1s-1s σ overlap
- HF: 1s-2p σ overlap
- O₂: One σ bond (pₓ-pₓ) + One π bond (p₂-p₂ parallel overlap)
σ = Straight/simple bond (head-to-head), π = Extra bond in double/triple bonds (side-to-side). Single bond = 1σ, Double bond = 1σ+1π, Triple bond = 1σ+2π!
Hybridization
Definition: Process of mixing atomic orbitals of different energies/shapes to form new equivalent orbitals of same energy/shape.
| Particular | sp³ | sp² | sp |
|---|---|---|---|
| Definition | Mixing of one s and three p orbitals | Mixing of one s and two p orbitals | Mixing of one s and one p orbital |
| Geometry | Tetrahedral (109.5°) | Trigonal planar (120°) | Linear (180°) |
| % s & p character | 25% s, 75% p | 33.3% s, 66.7% p | 50% s, 50% p |
| Examples | CH₄, NH₃, H₂O | BF₃, C₂H₄ | BeCl₂, C₂H₂ |
- Hybridization occurs in single atom only
- Only orbitals of comparable energies can mix
- Number of hybrid orbitals = Number of orbitals mixed
- Greater s-character = Shorter and stronger bonds
Hybridization: Count electron pairs around central atom → 4 pairs = sp³ (tetrahedral), 3 pairs = sp² (trigonal), 2 pairs = sp (linear)!
Bond Energy & Bond Length
Bond Energy: Average energy required to break all bonds of particular type in one mole of substance (kJ/mol).
Bond Length: Distance between nuclei of two atoms forming covalent bond.
| Bond | Bond Energy (kJ/mol) | Bond | Bond Length (pm) |
|---|---|---|---|
| C—C | 348 | C—C (sp³) | 154 |
| C=C | 614 | C=C (sp²) | 133 |
| C≡C | 839 | C≡C (sp) | 120 |
| H—H | 436 | B—F (sp²) | 130 |
| O=O | 495 | C=O (sp²) | 122 |
- Electronegativity difference of bonded atoms
- Size of atoms (larger = weaker bond)
- Bond length (shorter = stronger bond)
- Nature of orbital (more s-character = stronger)
- Bond order (higher = stronger)
Bond energy ∝ 1/Bond length. Triple bonds > Double bonds > Single bonds in strength. More s-character = Shorter, stronger bonds!
Dipole Moment
Definition: Product of electric charge and distance between opposite charged centres (µ = q × r).
1D = 3.336 × 10⁻³⁰ mC
Applications:
- Determine polarity: µ = 0 → non-polar, µ ≠ 0 → polar
- Calculate % ionic character: % = (µobserved/µtheoretical) × 100
- Stereochemistry: Distinguish cis-trans isomers
- Molecular geometry: Predict shape from dipole moment
- ∆E.N ≤ 0.4 → Non-polar covalent
- 0.4 < ∆E.N < 1.7 → Polar covalent
- ∆E.N = 1.7 → 50% ionic, 50% covalent
- ∆E.N > 1.7 → Ionic bond
Exception: H—F has ∆E.N = 1.9 but is polar covalent
Dipole moment = Measure of polarity. Symmetric molecules (CO₂, CCl₄) = Zero µ, Asymmetric molecules (H₂O, NH₃) = Non-zero µ!
Comparison of Ionic vs Covalent Compounds
| Property | Ionic Compounds | Covalent Compounds |
|---|---|---|
| Conductivity | Non-conductor (solid), Conductor (fused/aqueous) | Non-conductor (electrolytic solutions conduct) |
| Melting/Boiling Point | High | Low |
| Alignment | Non-directional, non-rigid | Directional, rigid |
| Crystalline properties | Show polymorphism & isomorphism | — |
| Reaction rate | Very high | Low |
| Solubility | Soluble in polar solvents | Like dissolves like |
| Isomerism | Do not show | Show isomerism |
Ionic = High MP/BP + Conducts when melted + Polar solvent soluble. Covalent = Low MP/BP + Poor conductors + Like dissolves like!
Applications & Importance
Practical Applications of Chemical Bonding Knowledge:
- Drug Design: Understanding bond strengths for effective medications
- Materials Science: Designing polymers, composites, alloys
- Semiconductor Industry: Silicon-germanium bonds in electronics
- Catalysis: Transition metal complexes with coordinate bonds
- Environmental Chemistry: Understanding pollutant binding
- Biochemistry: Protein folding, enzyme-substrate interactions
- Nanotechnology: Carbon nanotubes, graphene with sp² hybridization
- Energy Storage: Lithium-ion batteries with ionic interactions
Everything around you is held together by chemical bonds! Metals (metallic), water (polar covalent), salt (ionic), plastics (covalent) – bonding explains it all!