ISC Class 11 Chemistry Syllabus

Define atomic number (Z) and mass number (A).

Atomic number Z = number of protons in the nucleus (= electrons in a neutral atom). Mass number A = number of protons + number of neutrons.

What are isotopes? Give an example.

Isotopes are atoms of the same element with the same atomic number but different mass numbers (different number of neutrons). Example: protium, deuterium, tritium of hydrogen.

What are isobars? Give an example.

Isobars are atoms of different elements having the same mass number but different atomic numbers. Example: Ca-40 and Ar-40.

State the Modern Periodic Law.

The physical and chemical properties of elements are periodic functions of their atomic numbers. Given by Henry Moseley.

On what basis did Mendeleev arrange his periodic table, and what was its key defect?

Mendeleev arranged elements in increasing order of atomic mass. Its defect was anomalous pairs (e.g. Ar before K) and no fixed place for isotopes; resolved by using atomic number.

How many periods and groups are in the modern long-form periodic table?

7 periods (horizontal rows) and 18 groups (vertical columns).

Into which four blocks is the periodic table divided?

s-block, p-block, d-block (transition elements), and f-block (lanthanoids and actinoids), based on the subshell receiving the last electron.

Define atomic radius and how it varies across a period and down a group.

Atomic radius is the distance from the nucleus to the outermost shell. It decreases across a period (left to right) and increases down a group.

Define ionization enthalpy and its periodic trend.

The energy required to remove the most loosely bound electron from an isolated gaseous atom. It increases across a period and decreases down a group.

Define electron gain enthalpy and its trend.

The enthalpy change when an electron is added to a neutral gaseous atom. It becomes more negative across a period and less negative down a group (halogens have the most negative).

Define electronegativity and its periodic trend.

The tendency of an atom to attract the shared pair of electrons in a covalent bond. It increases across a period and decreases down a group; fluorine is most electronegative.

Why is the second ionization enthalpy always greater than the first?

Removing an electron from a positively charged cation requires more energy because the remaining electrons are held more tightly by the same nuclear charge.

What are valence electrons?

The electrons present in the outermost shell of an atom, which take part in chemical bonding and determine the valency of the element.

State the octet rule.

Atoms tend to gain, lose, or share electrons to achieve a stable configuration of eight electrons in their valence shell (a noble gas configuration).

Define ionic (electrovalent) bond.

A bond formed by the complete transfer of one or more electrons from one atom to another, resulting in oppositely charged ions held together by electrostatic attraction. Example: NaCl.

What conditions favour the formation of an ionic bond?

Low ionization enthalpy of the metal, high (negative) electron gain enthalpy of the non-metal, high lattice enthalpy, and a large electronegativity difference between the atoms.

Define covalent bond.

A bond formed by the mutual sharing of one or more pairs of electrons between two atoms. Example: H2, Cl2, CH4.

Distinguish between single, double, and triple covalent bonds.

Single bond shares one electron pair (1 sigma), double bond shares two pairs (1 sigma + 1 pi), triple bond shares three pairs (1 sigma + 2 pi). Bond strength increases and length decreases from single to triple.

Define bond length and bond enthalpy.

Bond length is the equilibrium distance between nuclei of two bonded atoms. Bond enthalpy is the energy required to break one mole of a particular bond in the gaseous state.

Define bond order and bond angle.

Bond order is the number of bonds between two atoms in a molecule. Bond angle is the angle between two adjacent bonds at the central atom.

What is a coordinate (dative) covalent bond?

A covalent bond in which the shared pair of electrons is contributed by only one of the two atoms. Example: NH4+ and the bond in O3.

What is a Lewis (electron dot) structure?

A representation of a molecule showing the bonding pairs and lone pairs of valence electrons as dots/lines around the symbols of atoms.

What is formal charge and how is it calculated?

Formal charge = (valence electrons) - (non-bonding electrons) - 1/2(bonding electrons). It helps identify the most stable Lewis structure.

What makes a covalent bond polar?

A difference in electronegativity between the bonded atoms causes unequal sharing of electrons, creating partial positive and negative charges. Example: HCl.

Define dipole moment and give its unit.

Dipole moment (mu) = charge x distance of separation; it measures bond/molecular polarity. Unit is the Debye (D); it is a vector quantity.

Why is CO2 nonpolar despite having polar bonds?

CO2 is linear and symmetrical, so the two equal and opposite C=O bond dipoles cancel out, giving a net dipole moment of zero.

What is the main postulate of VSEPR theory?

Electron pairs (bonding and lone) around a central atom repel each other and arrange themselves to minimize repulsion, giving the molecule a definite shape.

What is the order of repulsion strength between electron pairs in VSEPR theory?

Lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.

Using VSEPR, give the shapes of CH4, NH3, and H2O and their bond angles.

CH4 is tetrahedral (109.5 degrees); NH3 is trigonal pyramidal (107 degrees); H2O is bent/angular (104.5 degrees). Lone pairs reduce the angle.

What is hybridization?

The intermixing of atomic orbitals of slightly different energies to form an equal number of new, equivalent hybrid orbitals with definite geometry.

Give the geometry associated with sp, sp2, and sp3 hybridization.

sp is linear (180 degrees), sp2 is trigonal planar (120 degrees), and sp3 is tetrahedral (109.5 degrees).

What are the three states of matter and how do they differ in intermolecular forces?

Solid (strong forces, fixed shape and volume), liquid (moderate forces, fixed volume but no fixed shape), gas (negligible forces, no fixed shape or volume).

Define surface tension and viscosity (properties of the liquid state).

Surface tension is the force per unit length acting on the surface of a liquid, minimizing surface area. Viscosity is the resistance to the flow of a liquid due to internal friction between layers.

How do surface tension and viscosity vary with temperature?

Both surface tension and viscosity decrease with an increase in temperature, because higher kinetic energy weakens intermolecular attractions.

What is vapour pressure and how does it relate to boiling point?

Vapour pressure is the pressure exerted by a vapour in equilibrium with its liquid at a given temperature. A liquid boils when its vapour pressure equals the external (atmospheric) pressure.

List the main types of bonding/intermolecular forces.

Ionic bond, covalent bond, metallic bond, and intermolecular forces such as London dispersion forces, dipole-dipole interactions, and hydrogen bonding.