Class 11 Chemistry: CHEMICAL BONDING AND MOLECULAR STRUCTURE (Only 4.1)

4. CHEMICAL BONDING AND MOLECULAR STRUCTURE

4.1 KÖssel–Lewis Approach to Chemical Bonding
4.1.1 Octet Rule
4.1.2 Covalent Bond
4.1.3 Lewis Representation of Simple Molecules (the Lewis Structures)
4.1.4 Formal Charge
4.1.5 Limitations of the Octet Rule

When NaCl dissolves in water, it splits into Na⁺ and Cl⁻ ions that move freely, conducting electricity. The ions — not free electrons — carry charge through the solution. Solid NaCl cannot conduct because its ions are locked in the crystal lattice; only when they dissociate in water do they act as charge carriers.

Sodium metal is highly reactive because it readily loses electrons to achieve a stable noble-gas configuration; that rapid oxidation in water releases lots of energy and hydrogen gas. In NaCl, sodium exists as Na⁺ and chlorine as Cl⁻ — both already have electron configurations analogous to noble gases, so the ionic compound is stable and non-reactive under normal conditions.

Magnesium has two valence electrons in its outer shell and can lose both to form Mg²⁺, reaching a noble-gas-like configuration. Each chlorine atom gains one electron to become Cl⁻, so two chlorine atoms balance one Mg²⁺, giving the stable ionic formula MgCl₂. This electron transfer is energetically favorable and explains the product formed.

Ionic solids like NaCl are held together by strong electrostatic attraction between oppositely charged ions. Under stress, ionic layers can shift so like charges align; the resulting strong repulsive forces cause the crystal to cleave and shatter. That combination of strong bonding and rigid lattice yields hardness but also brittleness.

Noble gases such as helium have complete valence shells (a full duet for helium, full octets for larger noble gases), providing maximal electronic stability. Because their outer shells are already filled they have little tendency to gain, lose, or share electrons — making them chemically inert and safe for uses like lifting balloons.

Propane and butane molecules are held by strong covalent bonds; each atom satisfies its octet (or duet for hydrogen), so there are no unpaired electrons or free ions to react spontaneously at room temperature. Energy input (a spark or flame) is needed to break bonds and start combustion.

Fluorine atoms are very small and highly electronegative, strongly attracting electrons to complete their octet; as a result elemental F₂ is extremely reactive and forms bonds readily (often violently). In toothpaste we use stable fluoride ions (F⁻), not elemental fluorine, which are safe and beneficial in small amounts.

Rusting is an oxidation process: iron atoms lose electrons (are oxidized) to oxygen in the presence of moisture, forming iron oxides and hydroxides. By losing electrons, the iron atoms move toward a more stable electronic arrangement — the whole electrochemical process explains the familiar corrosion of metal structures.

Pure water contains very few ions and therefore cannot carry current well. Adding salt (NaCl) produces Na⁺ and Cl⁻ ions in solution that are free to move under an electric field — these mobile ions are the charge carriers that make the tester glow.

Plastics are polymers: long chains of repeating units bonded by strong covalent bonds. These covalent networks resist breakage and distribute stress across the chain, giving materials high strength-to-weight ratios. Chemical resistance also arises from the stability of the covalent backbone.

Molecular oxygen (O₂) contains a double covalent bond where each oxygen shares two pairs of electrons, completing their octets. This shared electron arrangement is energetically favorable and explains why O₂ is a stable, diatomic gas under normal conditions.

Carbon is tetravalent — it can form four covalent bonds — and it readily bonds to other carbon atoms (catenation), creating long chains, branched frameworks, and rings. This versatility is the fundamental reason for the vast diversity of organic compounds found in fuels, polymers, and living systems.

Carbon dioxide (CO₂) is linear with carbon forming two double bonds to oxygen. Each atom attains an octet through these shared pairs, producing a stable, symmetric, and relatively non-reactive molecule under ordinary atmospheric conditions.

Ammonia (NH₃) has a lone pair on nitrogen and a trigonal pyramidal shape, creating a permanent dipole that interacts favorably with polar water molecules. Methane (CH₄) is tetrahedral and symmetric, producing no net dipole; its non-polar nature prevents it from dissolving appreciably in water.

Ozone (O₃) has resonance structures with electrons delocalized across the molecule. These electronic arrangements allow ozone to absorb ultraviolet radiation, converting the energy into chemical forms and thereby protecting life on Earth while still being reactive enough to reform the ozone layer.

Solid CO₂ (dry ice) is composed of neutral covalent CO₂ molecules held together by weak van der Waals forces. There are no mobile ions or delocalized electrons in the solid state, so it cannot conduct electricity like ionic solids or metals.

Nitrate is a resonance-stabilized ion: the negative charge is delocalized over three equivalent oxygen atoms. Formal charge calculations and resonance structures together show that the negative charge is shared, which explains uniform reactivity and bond lengths observed experimentally.

Formal charge analysis helps minimize charge separation and select the Lewis structure that best reflects electron distribution. The structure with the smallest and most appropriate formal charges usually corresponds to lower energy and higher stability — a practical tool in predicting molecular arrangements.

The carbonate ion features resonance that delocalizes the two negative charges across three equivalent oxygen atoms. This delocalization lowers the ion’s energy and contributes to the stability of carbonate salts like chalk, making them relatively resistant to decomposition under mild conditions.

Elements in the third period, such as phosphorus, have access to vacant d-orbitals which allow them to expand their valence shell beyond eight electrons. This capacity to accommodate more electrons explains stable molecules with expanded octets like PCl₅ and SF₆.