CBSE Class 11 Chemistry

Chapters: 10

1. Some Basic Concepts of Chemistry

Matter and its measurement · Laws of chemical combination and Dalton's theory · Mole concept · Stoichiometry

Matter and Its Measurement – Matter is anything that has mass and takes up space. It can be solid, liquid or gas, and it can be an element, a compound or a mixture. Chemists measure matter in SI units. Every measurement has some doubt, so we write it with the right number of significant figures, use scientific notation for very big or small numbers, and change units with conversion factors.
Laws of Chemical Combination and Dalton's Atomic Theory – Five laws describe how elements combine. Mass is conserved in a reaction. A compound always has the same elements in the same mass ratio. When two elements make more than one compound, the masses of one that join a fixed mass of the other are in a small whole-number ratio. Reacting gases have volumes in simple whole-number ratios, and equal volumes of gases hold equal numbers of molecules. Dalton's atomic theory explains all of these with atoms.
Mole Concept, Molar Mass and Chemical Formulas – Atoms are far too small to count one by one, so chemists count them by weighing. Atomic masses are given in u, where 1 u is one-twelfth the mass of a carbon-12 atom. One mole is 6.022 × 10²³ particles, and its mass in grams equals the formula mass in u. With moles we can find the percentage of each element in a compound, and work back from percentages to the empirical and molecular formulas.
Stoichiometry, Limiting Reagent and Concentration Terms – Stoichiometry means measuring the amounts of substances in a reaction. A balanced equation acts like a recipe: its numbers give the mole ratio of reactants and products. With it you can change any mass into moles, use the ratio, and change back to the mass of any other substance. The reactant that runs out first is the limiting reagent; it decides how much product forms. For solutions, concentration is given as mass percent, mole fraction, molarity or molality.

2. Structure of Atom

Subatomic particles and early models · Towards Bohr's model · Quantum mechanical model

Subatomic Particles and Early Atomic Models – An atom is made of three main particles. Electrons (charge −1, very light) were found with cathode ray tubes. Protons (charge +1, about 1836 times heavier) were found in canal rays. Neutrons (no charge, mass about a proton) were found by Chadwick in 1932. Thomson pictured the atom as a ball of positive charge with electrons stuck in it. Rutherford fired alpha particles at gold foil and found a tiny, heavy, positive nucleus with electrons moving outside. Atomic number Z = number of protons; mass number A = protons + neutrons. Rutherford's model could not explain why atoms are stable or why they give line spectra, which led to Bohr's model.
Towards Bohr's Model: Light, Photons, Spectra and Bohr's Atom – Light is an electromagnetic wave: c = νλ, and wavenumber ν̄ = 1/λ. But some facts need particles: Planck said energy comes in packets (quanta) E = hν. Einstein used photons to explain the photoelectric effect: an electron comes out only if hν is above the work function W₀ = hν₀, and its kinetic energy is hν − hν₀. Atoms give line spectra, which means electron energies are fixed. Bohr put the electron on fixed orbits with angular momentum nh/2π; energy Eₙ = −2.18 × 10⁻¹⁸ Z²/n² J, radius rₙ = 52.9 n²/Z pm. A jump between orbits gives a photon, and 1/λ = R_H(1/n₁² − 1/n₂²) gives the Lyman, Balmer, Paschen, Brackett and Pfund series. Bohr works only for one-electron species.
Quantum Mechanical Model: Quantum Numbers, Orbitals and Configuration – Moving electrons also behave like waves: λ = h/mv (de Broglie). Because of this we cannot know exact position and momentum together: Δx·Δp ≥ h/4π (Heisenberg). So we drop sharp orbits and use orbitals: regions where the electron is likely to be, from the Schrödinger equation (ψ² gives probability). Four quantum numbers describe each electron: n (shell, size, energy), l (subshell, shape: 0 to n−1), mₗ (orientation: −l to +l) and mₛ (spin: +½ or −½). s is spherical, p is dumbbell, d is mostly four-lobed. Electrons fill by the Aufbau (n + l) rule, Pauli principle (max 2, opposite spins) and Hund's rule (single first). Half-filled and fully filled subshells are extra stable, so Cr is 3d⁵4s¹ and Cu is 3d¹⁰4s¹.

3. Classification of Elements and Periodicity in Properties

Periodic classification · Configuration and blocks · Periodic trends

Periodic Classification of Elements – Scientists sorted elements so that similar ones sit together. Döbereiner made triads, Newlands found that every 8th element repeats (octaves), Mendeleev arranged elements by atomic mass and left gaps for unknown ones. Moseley showed that atomic number is the real key. Modern periodic law: the properties of elements are a periodic function of their atomic numbers. The modern table has 7 periods and 18 groups. Elements with Z > 100 get temporary IUPAC names built from digit roots (nil, un, bi, tri, quad, pent, hex, sept, oct, enn) plus -ium.
Electronic Configuration and the s, p, d, f Blocks – The periodic table is really a map of electron filling. The period number is the highest shell (n) being used. The block is the subshell that receives the last electron: s-block (groups 1–2, ns¹⁻²), p-block (groups 13–18, ns² np¹⁻⁶), d-block (groups 3–12, (n−1)d¹⁻¹⁰ ns⁰⁻²) and f-block (lanthanoids and actinoids, (n−2)f¹⁻¹⁴ (n−1)d⁰⁻¹ ns²). Group: s-block = number of ns electrons; p-block = 10 + ns + np electrons; d-block = ns + (n−1)d electrons.
Periodic Trends in Properties – Two forces decide almost every trend: the pull of the nucleus (effective nuclear charge) and the distance of the outer shell. Across a period the nuclear pull grows while the shell stays the same, so atoms shrink, ionisation enthalpy rises, electron gain enthalpy becomes more negative and electronegativity rises. Down a group a new shell is added, so atoms grow and these values fall. Cations are smaller and anions bigger than their atoms. Valence follows the outer electrons; metallic reactivity is highest at the bottom left and non-metallic reactivity at the top right.

4. Chemical Bonding and Molecular Structure

Kossel-Lewis approach and ionic bond · Bond parameters · VSEPR theory · Valence bond theory and hybridisation · Molecular orbital theory · Hydrogen bonding

Chemical Bonding: Ionic, Covalent and Metallic Bonds – Atoms join together (bond) to become more stable. Only their outer electrons take part. Most atoms are most stable with 8 outer electrons: the octet rule. There are three main ways to reach it. In an ionic bond, a metal gives electrons to a non-metal, making oppositely charged ions that attract. In a covalent bond, two non-metals share pairs of electrons to make molecules. In a metallic bond, metal atoms release outer electrons into a shared 'sea' that holds positive ions together. The difference in electronegativity (how strongly an atom pulls shared electrons) tells us which kind of bond forms. The type of bond explains melting points, whether a substance conducts electricity, and whether it dissolves in water.
Kossel-Lewis Approach and the Ionic Bond – Atoms join so that each gets a stable outer shell of 8 electrons (an octet), like a noble gas. Kossel said atoms can give or take electrons to make ions (ionic bond). Lewis said atoms can also share pairs of electrons (covalent bond). Lewis structures show these electrons as dots and lines. Formal charge (V − L − B/2) helps pick the best Lewis structure. Ions pack into a crystal, and the energy released is linked to the lattice enthalpy.
Bond Parameters, Resonance and Polarity – A covalent bond is described by four numbers: bond length (distance between nuclei), bond angle (angle between bonds at an atom), bond enthalpy (energy to break 1 mol of bonds) and bond order (number of shared pairs). Higher bond order means a shorter, stronger bond. When one Lewis structure cannot describe a molecule, the real molecule is a resonance hybrid of several structures. Unequal sharing makes a bond polar; the dipole moment μ = q × d measures it, and the shape decides whether bond dipoles cancel.
VSEPR Theory: Shapes of Molecules – VSEPR stands for Valence Shell Electron Pair Repulsion. The electron pairs around a central atom repel each other and move as far apart as possible. The number of pairs sets the basic arrangement: 2 linear, 3 trigonal planar, 4 tetrahedral, 5 trigonal bipyramidal, 6 octahedral. Lone pairs take more room than bond pairs (lp–lp > lp–bp > bp–bp), so they squeeze bond angles and change the shape we see, for example NH₃ (pyramidal, 107°) and H₂O (bent, 104.5°).
Valence Bond Theory and Hybridisation – Valence bond theory says a covalent bond forms when half-filled orbitals of two atoms overlap and the two electrons pair up with opposite spins. Head-on overlap makes a strong σ (sigma) bond; side-on overlap of p orbitals makes a weaker π (pi) bond. To explain real shapes, an atom first mixes its orbitals into equal hybrid orbitals: sp (linear, 180°), sp² (trigonal planar, 120°), sp³ (tetrahedral, 109.5°), sp³d (trigonal bipyramidal) and sp³d² (octahedral).
Molecular Orbital Theory – In molecular orbital (MO) theory, atomic orbitals of the bonded atoms combine into molecular orbitals that belong to the whole molecule. By LCAO, adding two atomic orbitals gives a lower-energy bonding MO; subtracting gives a higher-energy antibonding MO. Electrons fill MOs by the Aufbau, Pauli and Hund rules. Bond order = ½ (Nb − Na). A positive bond order means the molecule exists; unpaired electrons make it paramagnetic. MO theory explains why He₂ does not exist and why O₂ is paramagnetic.
Hydrogen Bonding – A hydrogen bond is a weak attraction between an H atom that is joined to a small, very electronegative atom (F, O or N) and a lone pair on another such atom. It is shown with a dotted line: X–H···Y. It is much weaker than a covalent bond (about 10–40 kJ mol⁻¹) but strong enough to raise boiling points, make ice float and hold DNA together. Intermolecular H-bonds join different molecules; intramolecular H-bonds form inside one molecule.

5. Chemical Thermodynamics

Thermodynamic terms · Measuring ΔU and ΔH · Enthalpies of different reactions · Spontaneity and Gibbs energy

First Law of Thermodynamics: System, Heat, Work and ΔU – Thermodynamics tracks energy. The part we study is the system; the rest is the surroundings; the wall between them is the boundary. A system can be open, closed or isolated. Internal energy U is the total energy stored in the system. It changes only in two ways: by heat q or by work w. The first law says ΔU = q + w (IUPAC signs: q and w are positive when energy goes INTO the system). U is a state function; q and w are path functions. For expansion against a constant outside pressure, w = −p_ext ΔV; for a reversible isothermal expansion of an ideal gas, w = −2.303 nRT log(V₂/V₁).
First Law of Thermodynamics – Internal energy U is the total energy of the molecules inside a system. It changes in only two ways: by heat Q (energy that flows because of a temperature difference) and by work W (energy moved by a force, like a moving piston). First law: ΔQ = ΔU + ΔW. Heat given to a gas partly raises its internal energy and partly lets it do work. Work by a gas at constant pressure is PΔV. For an ideal gas Cp − Cv = R.
Enthalpy, Calorimetry and Hess's Law – Enthalpy H = U + pV. At constant pressure the heat taken in or given out equals ΔH; at constant volume it equals ΔU. We measure ΔU in a sealed bomb calorimeter and ΔH in an open cup calorimeter, using q = C ΔT or q = m c ΔT. For gases, ΔH = ΔU + Δn_g RT. The reaction enthalpy ΔrH is negative for exothermic and positive for endothermic reactions, and ΔrH° = ΣΔfH°(products) − ΣΔfH°(reactants). Hess's law: the total enthalpy change is the same whether a reaction happens in one step or many, so thermochemical equations can be added like algebra.
Enthalpies of Different Types of Reactions – Each kind of change has its own named enthalpy, always for 1 mol. Phase changes: fusion (melting), vaporisation (boiling) and sublimation (ΔsubH = ΔfusH + ΔvapH); all are endothermic. Combustion: 1 mol burns completely in oxygen; always exothermic. Atomisation: 1 mol of a substance breaks fully into gaseous atoms. Bond dissociation enthalpy: 1 mol of a given bond is broken in the gas phase; for polyatomic molecules we use mean bond enthalpy. ΔrH ≈ Σ(bonds broken) − Σ(bonds formed). Lattice enthalpy: 1 mol of an ionic solid is split into gaseous ions (found by a Born–Haber cycle). Enthalpy of solution = lattice enthalpy + hydration enthalpy.
Spontaneity, Entropy and Gibbs Energy – A spontaneous process happens by itself, without outside help (it may be fast or slow). ΔH alone cannot predict it: ice melts on its own although it takes in heat. Entropy S measures how spread out energy and matter are; ΔS = q_rev/T. Second law: for any spontaneous change, ΔS_total = ΔS_system + ΔS_surroundings > 0; at equilibrium it is zero. Third law: a perfect crystal at 0 K has zero entropy. Gibbs energy G = H − TS combines both: at constant T and p, ΔG = ΔH − TΔS, and ΔG < 0 means spontaneous, ΔG = 0 means equilibrium. It links to the equilibrium constant by ΔG° = −RT ln K = −2.303 RT log K.

6. Equilibrium

Equilibrium in physical and chemical processes · Ionic equilibrium

Chemical Equilibrium: Kc, Kp, Q and Gibbs Energy – In a closed container a reversible reaction goes both ways. After some time the forward and backward rates become equal, so amounts stop changing, but the reaction does not stop. This is dynamic equilibrium. At equilibrium the ratio of products to reactants (each raised to its coefficient) is a fixed number, the equilibrium constant K. Kc uses concentrations, Kp uses partial pressures, and Kp = Kc(RT)^Δn. Pure solids and liquids are left out of K. The reaction quotient Q tells the direction: Q < K goes forward, Q > K goes backward, Q = K is equilibrium. K and Gibbs energy are linked: ΔG = ΔG° + RT ln Q and ΔG° = −RT ln K.
Le Chatelier's Principle: Factors Affecting Equilibrium – Le Chatelier's principle says: if you disturb a system at equilibrium, it shifts in the direction that reduces the disturbance. Add a reactant or remove a product → shifts forward. Increase pressure (smaller volume) → shifts to the side with fewer gas moles. Raise temperature → shifts in the heat-absorbing (endothermic) direction, and K changes. A catalyst only helps reach equilibrium faster; it does not move it. Inert gas at constant volume does nothing; at constant pressure it acts like lowering pressure. Only temperature changes the value of K.
Buffer Solutions and Solubility Product – A buffer is a solution that keeps its pH nearly the same when a little acid or base is added. An acidic buffer is a weak acid with its salt (CH₃COOH + CH₃COONa); a basic buffer is a weak base with its salt (NH₄OH + NH₄Cl). Its pH is given by the Henderson equation: pH = pKa + log([salt]/[acid]). A sparingly soluble salt in its saturated solution sets up an equilibrium with its ions; the product of ion concentrations (each raised to its coefficient) is the solubility product Ksp. For AB, Ksp = s²; for AB₂ or A₂B, Ksp = 4s³. If the ionic product Q exceeds Ksp, a precipitate forms. A common ion lowers solubility.
Ionic Equilibrium: Acids, Bases, Ionisation and pH – Acids and bases can be defined in three ways: Arrhenius (give H⁺ or OH⁻ in water), Brønsted–Lowry (proton donor or acceptor) and Lewis (electron-pair acceptor or donor). Strong acids and bases ionise fully; weak ones ionise only a little and set up an equilibrium with constant Ka or Kb. Water itself ionises: Kw = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25 °C. pH = −log[H⁺] and pH + pOH = 14. For a weak electrolyte, α ≈ √(K/C), so dilution increases α (Ostwald's dilution law). For a conjugate pair, Ka × Kb = Kw. Salts of weak acids or weak bases react with water (hydrolysis) and give non-neutral solutions.

7. Redox Reactions

Redox reactions

Balancing Redox Reactions: Oxidation Number and Half-Reaction Methods – In a balanced redox equation, three things must match: atoms, charge, and electrons. The electrons lost by the reducing agent must equal the electrons gained by the oxidising agent. The oxidation number method balances the change in numbers first; the half-reaction method splits the reaction into two halves and joins them. O is balanced with H₂O, and H with H⁺ in acid or with OH⁻ in base.
Oxidation Number: Rules, Calculation and Types of Redox – Oxidation number is a pretend charge. We imagine that shared electrons go fully to the atom that pulls harder. A few simple rules give the number for any atom, and all numbers in a species add up to its charge. If an atom’s oxidation number goes up, it is oxidised; if it goes down, it is reduced. With this we can sort redox reactions into combination, decomposition, displacement and disproportionation.
Redox Reactions and Electrode Processes – If we keep the oxidation half and the reduction half in separate beakers and join them with a wire and a salt bridge, the electrons travel through the wire. This makes a cell. Each metal in its ion solution is an electrode with its own potential. The hydrogen electrode is set at 0.00 V. Listing electrodes by standard potential gives the electrochemical series, and E°cell = E°(cathode) − E°(anode).
Redox Reactions: From Oxygen to Electron Transfer – Oxidation first meant adding oxygen or removing hydrogen. Reduction meant the opposite. Today we use a bigger idea: oxidation is losing electrons and reduction is gaining electrons. Both always happen together, so we call them redox reactions. A more active metal gives electrons to the ion of a less active metal.

8. Organic Chemistry: Some Basic Principles and Techniques

Structure and representation · Nomenclature and isomerism · Reaction mechanism basics · Purification and analysis

Organic Compounds: Structure, Formulae and Classification – Organic chemistry is the chemistry of carbon compounds. A carbon atom has 4 outer electrons, so it always makes 4 bonds. This is called tetravalence. With 4 single bonds the carbon is sp³ and its bonds point to the corners of a tetrahedron (109.5°). With one double bond it is sp² and flat (120°). With a triple bond it is sp and straight (180°). We can draw the same molecule in several ways: complete, condensed, bond-line and 3D wedge-dash. Finally we sort organic compounds into open-chain and ring compounds, and into families by their functional group. Members of one family form a homologous series, where each member differs by one CH₂.
IUPAC Nomenclature and Isomerism of Organic Compounds – Millions of organic compounds exist, so every one needs a clear, unique name. The IUPAC system builds a name from three parts: the word root (number of carbons in the main chain), the suffix (type of bond and main functional group) and the prefixes (side groups). You pick the longest chain that holds the main group, number it so the main group and branches get the lowest numbers, and write side groups in alphabetical order. Isomers are compounds with the same molecular formula but different arrangement. Structural isomers differ in how atoms are joined (chain, position, functional group, metamerism). Stereoisomers are joined the same way but differ in 3D arrangement (geometrical cis/trans and optical isomers).
Basics of Organic Reaction Mechanism – A reaction mechanism is the step-by-step story of which bonds break, which form, and where the electrons move. A covalent bond can break in two ways. In homolysis each atom takes one electron and free radicals form. In heterolysis one atom takes both electrons, giving a carbocation (C with + charge) and a carbanion or an anion. Electron-poor species that seek electrons are electrophiles; electron-rich species that give electrons are nucleophiles. The movement of electrons inside a molecule is controlled by four effects: the inductive effect (pull along σ bonds), the resonance effect (spreading of π electrons), the electromeric effect (a temporary full shift when a reagent attacks) and hyperconjugation (σ C–H electrons spreading into a nearby empty p orbital or π bond). These effects decide how stable an intermediate is and where a reagent attacks.
Purification and Analysis of Organic Compounds – A compound made in a lab or taken from a plant is never pure at first. We purify it by using a difference in properties: solubility (crystallisation), boiling point (distillation and its types), sublimation, solubility in two liquids (differential extraction) or how strongly it sticks to a surface (chromatography). A pure solid has a sharp melting point. Next we find which elements are present (qualitative analysis): carbon and hydrogen by heating with copper(II) oxide, and N, S, halogens and P by Lassaigne's sodium fusion test. Finally we find how much of each element is present (quantitative analysis): Liebig's method for C and H, Dumas and Kjeldahl methods for N, Carius method for halogens and S, and oxygen by difference. From the percentages we can get the empirical formula.

9. Hydrocarbons

Alkanes · Alkenes · Alkynes · Aromatic hydrocarbons

Alkanes – Alkanes are hydrocarbons with only single C–C bonds. Their general formula is CnH2n+2. We make them by adding hydrogen to alkenes, by the Wurtz reaction, by reducing alkyl halides and by decarboxylation. They are non-polar, burn with a lot of heat and swap H for Cl in light. Around the C–C bond the molecule can twist into staggered and eclipsed shapes called conformations.
Alkenes – Alkenes are hydrocarbons with one C=C double bond and the formula CnH2n. The double bond is one σ bond plus one π bond, so the molecule is flat around it and cannot twist. That gives cis and trans isomers. We make alkenes by removing small molecules (H₂O, HX, X₂) or by partly adding H₂ to alkynes. Alkenes mainly do addition reactions: H₂, X₂, HX (Markovnikov rule, or anti-Markovnikov with peroxide), water, ozone, KMnO₄ and polymerisation.
Alkynes – Alkynes have a C≡C triple bond and the formula CnH2n−2. The two triple-bond carbons are sp hybridised, so the molecule is a straight line. A hydrogen on a triple-bond carbon is weakly acidic because an sp carbon holds electrons tightly: strong bases like NaNH₂ remove it. The triple bond can open twice, so alkynes add two molecules of H₂, X₂ or HX. With water (Hg²⁺ catalyst) they give aldehydes or ketones, and three ethyne molecules can join into benzene.
Benzene and Aromaticity – Benzene (C₆H₆) is a flat ring of six sp² carbons. Its six π electrons are spread over the whole ring, so all C–C bonds are equal and the ring is extra stable. A ring like this is called aromatic when it is cyclic, planar, fully conjugated and has 4n + 2 π electrons (Hückel rule). Benzene prefers electrophilic substitution (nitration, halogenation, sulphonation, Friedel–Crafts) over addition, because substitution keeps the stable ring. A group already on the ring decides where the next group goes: ortho/para or meta. Benzene and fused-ring hydrocarbons are toxic and can cause cancer.

10. Formative-only topics

s- and p-block elements · The gaseous state

s- and p-Block Elements: Configuration and Trends – s-block = groups 1 and 2 (outer configuration ns¹ or ns²); p-block = groups 13 to 18 (ns² np¹⁻⁶). Down a group, atoms get bigger and ionisation enthalpy falls; across a period, atoms get smaller and ionisation enthalpy rises overall. Small ions have large hydration enthalpy. The first element of each group behaves differently (small size, high charge density, no d orbitals) and often resembles the element diagonally below it.
The Ideal Gas Law: PV = nRT – Boyle (P ∝ 1/V), Charles (V ∝ T) and Avogadro (V ∝ n) combine into one rule: PV = nRT, with R = 8.314 J mol⁻¹ K⁻¹. Temperatures must be in kelvin. It also gives molar mass M = dRT/P. Real gases follow it well at low pressure and high temperature; at high pressure or low temperature they deviate, measured by Z = PV/nRT.