NCERT Summary: Summary of Chemistry - 2

Chemical Bonding

Atoms and Subatomic Particles

• Atoms are composed of three principal subatomic particles: protons, neutrons and electrons.
• Protons and neutrons are concentrated in the central nucleus. A proton carries a single positive charge; the number of protons in the nucleus is the atomic number and it uniquely identifies an element.
• Electrons carry a single negative charge and move in regions around the nucleus called orbitals or shells. In a neutral atom the number of electrons equals the number of protons.
• The chemical behaviour of an atom is largely determined by the electrons in its outermost shell (the valence shell), because these electrons participate in bond formation.

Electronic Shells and Stability

• Electrons occupy discrete shells; each shell can hold a limited number of electrons. A completely filled outer shell is a particularly stable arrangement (noble-gas configuration).
• Atoms tend to gain, lose or share electrons to achieve a stable electron configuration, commonly approximated by the octet rule (eight electrons in the valence shell) for many main-group elements.
• Valency (or valence) of an element can be viewed as the combining capacity of its atoms, often expressed as the number of hydrogen atoms with which one atom of the element can combine.

Nature of Chemical Bonding

Chemical bonds arise from electrostatic attractions that allow two or more atoms to be held together such that one or more electrons are simultaneously attracted by more than one nucleus. The principal types of bonding encountered in simple chemical compounds are described below.

Types of Chemical Bonds

• Ionic (Electrovalent) Bond: Formed when one atom transfers one or more electrons to another atom, producing positive and negative ions that attract each other. Ionic bonds are common between metals and non-metals. Characteristic features include high melting and boiling points, crystalline solids at room temperature, electrical conductivity in molten state or when dissolved in water, and generally good solubility in polar solvents such as water. Example: NaCl formed from Na and Cl.
• Covalent Bond: Formed when two atoms share one or more pairs of electrons. Covalent bonding is typical between non-metal atoms. Molecules with covalent bonds may be gases, liquids or solids; they often have lower melting and boiling points than ionic solids and do not conduct electricity in the neutral molecular form. Example: CH₄, H₂O.
• Coordinate (Dative) Covalent Bond: A special case of covalent bonding in which both electrons of a shared pair are contributed by the same atom (the donor), while the acceptor atom provides an empty orbital. The bond is represented in Lewis notation with an arrow from donor to acceptor. Coordinate bonds are still covalent in nature; the physical properties of a compound containing such bonds depend on its overall structure and cannot be generalised simply as "insoluble" or "non-conducting." Example: the ammonium ion NH⁴⁺ (formed when NH₃ donates a lone pair to H⁺).
• Metallic Bond: Present in metals, where valence electrons are delocalised over a lattice of positive ions. This "sea of electrons" explains metallic properties such as electrical and thermal conductivity, malleability and ductility, and metallic lustre. Example: metallic bonding in copper, iron.

Other Bonding Concepts and Models

• Lewis Structures and valence-shell electron pair repulsion (VSEPR) theory help predict the arrangement of electrons and the geometry of molecules.
• Polar covalent bonds occur when the shared electron pair is unequally distributed due to a difference in electronegativity; this produces a dipole moment in the bond. Ionic bonds can be considered an extreme case of polar covalency when electron transfer is effectively complete.
• Bond strength and bond length vary with bond order (single, double, triple bonds) and the atoms involved; higher bond order typically means shorter, stronger bonds.

Chemical Reactions and Equations

Basic Concepts

• Elements are substances composed of only one kind of atom. Examples: C (carbon), He (helium), Na (sodium), Fe (iron).
• Molecules are assemblies of two or more atoms bonded together. A molecule may contain atoms of the same element (e.g., O₂, O₃) or different elements (e.g., CH₄, NH₃).
• Compounds are substances made of two or more elements in fixed proportions, represented by a chemical formula (for example, NaCl, CH₄, H₂O).
• A chemical reaction (chemical change) is a process in which one or more substances (reactants) are converted into new substances (products). A chemical equation symbolically represents this change.

Chemical Equations and Conservation Laws

• A chemical equation shows reactants on the left and products on the right, usually separated by an arrow (→). Example: C + O₂ → CO₂.
• The Law of Conservation of Mass requires that the number of atoms of each element be the same on both sides of a chemical equation; therefore equations must be balanced.
• When a reaction can proceed in both directions and an equilibrium mixture results, it is denoted by a double-headed arrow (⇌). Example: N₂ + 3H₂ ⇌ 2NH₃.

Balancing a Simple Equation (Worked Example)

Balance the formation of water from hydrogen and oxygen:

Unbalanced: H₂ + O₂ → H₂O

Balance by changing the number of molecules (stoichiometric coefficients), not by altering formulas.

Place coefficient 2 before H2 on the reactant side and coefficient 2 before H₂O on the product side:

Balanced: 2H2 + O₂ → 2H₂O

This ensures there are four H atoms and two O atoms on each side, satisfying conservation of atoms.

Catalysts and Chemical Equilibrium

• A catalyst is a substance that increases the rate at which chemical equilibrium is attained by lowering the activation energy; it speeds both the forward and reverse reactions equally and therefore does not change the position of equilibrium.
• The position of equilibrium can be shifted by changing concentrations, pressure (for gaseous systems), or temperature; these shifts are qualitatively described by Le Chatelier's principle.

Classification of Chemical Reactions

• Combination (Synthesis) Reaction: Two or more substances combine to form a single product. Example: 2H₂ + O₂ → 2H₂O.
• Decomposition Reaction: A single substance breaks down into two or more simpler substances. Example: 2HgO → 2Hg + O₂ (on heating).
• Displacement (Single Displacement) Reaction: An element displaces another element from its compound. Example: Zn + CuSO₄ → ZnSO₄ + Cu.
• Double Displacement (Metathesis) Reaction: Exchange of ions between two compounds. Example: AgNO₃ + NaCl → AgCl (precipitate) + NaNO₃.
• Precipitation Reaction: A reaction that forms an insoluble solid (precipitate) when two aqueous solutions are mixed. Example: formation of AgCl.
• Oxidation-Reduction (Redox) Reaction: Reactions that involve transfer of electrons, or gain/loss of oxygen or hydrogen. Oxidation is loss of electrons (or gain of oxygen / loss of hydrogen); reduction is gain of electrons (or loss of oxygen / gain of hydrogen). The substance that causes oxidation and is itself reduced is called an oxidising agent; the one that causes reduction and is itself oxidised is a reducing agent. Example: 2Na + Cl₂ → 2NaCl (sodium is oxidised, chlorine is reduced).
• Exothermic and Endothermic Reactions: Exothermic reactions release heat (temperature of surroundings increases), while endothermic reactions absorb heat (temperature of surroundings decreases).

Key Principles

• Atoms are neither created nor destroyed in a chemical reaction; only the arrangements (bonds) of atoms change.
• Chemical reactions commonly proceed by making and breaking bonds between atoms.
• Industrial chemical processes frequently use catalysts and control of temperature, pressure and concentration to optimise yield and selectivity.

Matter and its Nature

Definition and Particle Nature of Matter

• Matter is anything that occupies space and has mass; it can be perceived through one or more senses and offers resistance to motion or force.
• Matter is composed of tiny particles (atoms, molecules or ions) that have space between them, are in constant motion and exert mutual attractive forces.

States of Matter

• Matter commonly exists in three classical states: solids, liquids and gases.
• Solids have a definite shape and volume, with particles closely packed in ordered arrangements and strong intermolecular attractions; solids are rigid and resist shape change.
• Liquids have definite volume but no fixed shape; they conform to the shape of their container. Particles are less tightly packed than in solids and can slide past each other, so liquids flow. Liquids are fluids.
• Gases have neither definite shape nor definite volume; gas particles are widely separated, move randomly and are highly compressible compared with solids and liquids.
• The magnitude of intermolecular forces is largest in solids, intermediate in liquids and smallest in gases; particle kinetic energy shows the opposite trend.

Interconversion of States and Heat Effects

• Changing temperature or pressure can convert matter from one state to another (melting, freezing, vaporisation, condensation, sublimation and deposition).
• Melting (Fusion) is the transition from solid to liquid. The temperature at which a solid melts at a given pressure is its melting point.
• Boiling is the formation of vapour bubbles within the liquid; the temperature at which a liquid boils at a given pressure is its boiling point.
• Latent heat of fusion is the heat required to convert 1 kg of a solid into liquid at its melting point (at constant pressure) without a change in temperature.
• Latent heat of vaporisation is the heat required to convert 1 kg of a liquid into vapour at its boiling point (at constant pressure) without a change in temperature.
• Sublimation is the direct transition between solid and gas phases without passing through the liquid state; the reverse process is called deposition. Example: iodine sublimes on heating; frost is formed by deposition of water vapour.

Evaporation and Cooling

• Evaporation is a surface phenomenon in which molecules at the liquid surface gain sufficient energy to enter the vapour state. Evaporation depends on surface area, temperature, humidity and air movement.
• Evaporation causes cooling because high-energy molecules leave the liquid, reducing the average kinetic energy and hence the temperature of the remaining liquid. This principle explains perspiration and the cooling effect of sweat.
• Condensation occurs when vapour loses energy on contact with a colder surface and converts to liquid, such as droplets forming on a cold glass.

Physical and Chemical Changes

• Physical properties (shape, size, colour, state, density, melting point, boiling point) are attributes that can be observed without changing the substance's chemical identity.
• A physical change alters only the physical form of a substance (for example, melting, freezing, dissolution) and is generally reversible; no new substance is formed.
• A chemical change (chemical reaction) alters the chemical composition, producing one or more new substances that usually have different properties from the reactants and are often not easily reversible.

Examples, Purity and Industrial Relevance

• Pure substances can often be obtained from mixtures or solutions by techniques such as crystallisation, distillation or filtration.
• Rusting of iron is a chemical change requiring both oxygen and moisture; the common product is hydrated iron(III) oxide (rust). A simplified net reaction can be represented as: 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃, which further dehydrates to form hydrated iron(III) oxide Fe₂O₃·xH₂O (rust).
• To protect iron from rusting, prevent contact with oxygen or water by painting, greasing, or by coating with a more reactive metal (for example, galvanisation deposits zinc on iron). Stainless steel contains iron alloyed with chromium, nickel and other elements to resist corrosion.

Energy Changes in Reactions

• Reactions that absorb heat from the surroundings are called endothermic; those that release heat are called exothermic. Whether a compound's formation from elements is endothermic or exothermic depends on bond energies and reaction pathway.

End of chapter summary: The particle view of matter explains states, changes of state and properties. Chemical bonding-ionic, covalent, coordinate and metallic-explains how atoms combine. Chemical reactions rearrange atoms and obey conservation of mass; they are classified by the type of change and often involve energy changes. Understanding these fundamentals is essential for further study in chemistry and for practical applications in industry, materials and daily life.