DPT Admission Test Chemistry Syllabus

Define an ionic bond and give a typical example.

An electrostatic attraction formed by complete transfer of electrons from a metal to a non-metal, e.g., NaCl (Na+ and Cl-).

Define a covalent bond and distinguish polar from non-polar covalent bonds.

A bond formed by sharing electron pairs; non-polar = equal sharing (similar electronegativity), polar = unequal sharing (electronegativity difference).

What is a coordinate (dative) covalent bond? Give an example.

A covalent bond in which both shared electrons come from one atom; e.g., NH4+ (N donates the lone pair to H+) or H3O+.

What approximate electronegativity difference distinguishes ionic, polar covalent, and non-polar covalent bonds?

Difference > ~1.7 = mostly ionic; ~0.4–1.7 = polar covalent; < ~0.4 = non-polar covalent.

What is hybridization, and what are the geometries of sp, sp2, and sp3 hybrids?

Mixing of atomic orbitals to form equivalent hybrid orbitals: sp = linear (180°), sp2 = trigonal planar (120°), sp3 = tetrahedral (109.5°).

What is the hybridization and shape of the carbon atom in methane (CH4)?

sp3 hybridized; tetrahedral shape with bond angles of 109.5°.

What is the hybridization, geometry, and bond angle of the carbon atoms in ethene (C2H4)?

sp2 hybridized, trigonal planar geometry, bond angle ~120°, with one sigma and one pi bond between the carbons.

Explain the bent shape and ~104.5° bond angle of the water molecule.

Oxygen is sp3 hybridized with two lone pairs; lone pair–bond pair repulsion compresses the H–O–H angle from 109.5° to ~104.5°.

According to VSEPR theory, what determines molecular geometry?

Electron pairs around the central atom arrange to minimize repulsion; lone pairs repel more strongly than bonding pairs, altering bond angles.

What is hydrogen bonding and which three elements bonded to hydrogen produce it?

A strong dipole–dipole attraction between H bonded to a highly electronegative atom (F, O, or N) and a lone pair on F, O, or N of another molecule.

List the intermolecular forces from weakest to strongest.

London dispersion forces < dipole–dipole forces < hydrogen bonding (all weaker than covalent/ionic bonds).

Why does water have an unusually high boiling point compared to H2S?

Water forms extensive hydrogen bonds (O is highly electronegative), whereas H2S has only weak dipole–dipole/dispersion forces.

What are London dispersion forces and which factor increases their strength?

Weak attractions from temporary instantaneous dipoles in all molecules; strength increases with molecular size/number of electrons (polarizability).

State Boyle's Law and give its equation.

At constant temperature, the volume of a fixed mass of gas is inversely proportional to its pressure: P1V1 = P2V2 (PV = constant).

State Charles's Law and give its equation.

At constant pressure, the volume of a fixed mass of gas is directly proportional to its absolute (Kelvin) temperature: V1/T1 = V2/T2.

Write the ideal gas equation and define each term.

PV = nRT, where P = pressure, V = volume, n = moles, R = gas constant (0.0821 L·atm·mol⁻¹·K⁻¹), and T = absolute temperature in Kelvin.

State Avogadro's Law.

Equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules; V ∝ n at constant T and P.

State Dalton's Law of Partial Pressures.

The total pressure of a gas mixture equals the sum of the partial pressures of its component gases: P_total = P1 + P2 + P3 + …

State Graham's Law of diffusion.

The rate of diffusion/effusion of a gas is inversely proportional to the square root of its molar mass: rate ∝ 1/√M.

Define vapor pressure and explain how it relates to a liquid's boiling point.

Pressure exerted by a vapor in equilibrium with its liquid; a liquid boils when its vapor pressure equals the external (atmospheric) pressure.

What is the difference between evaporation and boiling?

Evaporation is surface vaporization occurring at any temperature; boiling is bulk vaporization throughout the liquid when vapor pressure equals atmospheric pressure.

What is the difference between crystalline and amorphous solids?

Crystalline solids have a long-range ordered (regular) arrangement and sharp melting points; amorphous solids lack long-range order and soften over a range of temperatures.

Name the four main types of crystalline solids based on their constituent particles.

Ionic, covalent (network), molecular, and metallic solids.

What is a unit cell, and how many atoms are in a body-centered cubic (BCC) unit cell?

The smallest repeating structural unit of a crystal lattice; a BCC cell contains 2 atoms (8 corners × 1/8 + 1 center).

State the Law of Mass Action and write the equilibrium constant expression for aA + bB ⇌ cC + dD.

Reaction rate is proportional to the product of active masses (molar concentrations) of reactants. Kc = [C]^c[D]^d / [A]^a[B]^b.

State Le Chatelier's Principle.

If a system at equilibrium is disturbed by a change in concentration, pressure, or temperature, the equilibrium shifts to counteract that change.

For an exothermic reaction at equilibrium, how does increasing temperature affect the equilibrium position?

It shifts the equilibrium backward (toward reactants), decreasing the yield of products and lowering Kc.

Define pH and pOH, and state the relationship between them at 25°C.

pH = -log[H⁺] and pOH = -log[OH⁻]; at 25°C, pH + pOH = 14 (since Kw = [H⁺][OH⁻] = 1 × 10⁻¹⁴).

Define enthalpy of a reaction and state Hess's Law.

Enthalpy (ΔH) is the heat change at constant pressure; Hess's Law states the total enthalpy change of a reaction is independent of the pathway taken (depends only on initial and final states).

Define the rate of a reaction and list four main factors that affect it.

Change in concentration of reactant or product per unit time. Factors: concentration, temperature, surface area, and presence of a catalyst (also pressure for gases).