AKU MBBS Admission Test Chemistry Syllabus

What are the shapes and bond angles of water (H2O) and ammonia (NH3)?

H2O: bent/angular, ~104.5° (2 lone pairs). NH3: trigonal pyramidal, ~107° (1 lone pair). Both based on tetrahedral electron geometry.

Give the molecular geometries for 5 and 6 electron domains (no lone pairs).

5 domains: trigonal bipyramidal, angles 90° and 120°. 6 domains: octahedral, 90°.

What is hybridization?

The mixing of atomic orbitals of similar energy to form an equal number of new equivalent hybrid orbitals used in bonding, explaining observed molecular geometries.

Match hybridization (sp, sp2, sp3) to geometry and bond angle.

sp: linear, 180°. sp2: trigonal planar, 120°. sp3: tetrahedral, 109.5°.

State the hybridization of carbon in methane, ethene, and ethyne.

Methane (CH4): sp3. Ethene (C2H4): sp2. Ethyne (C2H2): sp.

Differentiate sigma (σ) and pi (π) bonds.

Sigma: head-on (axial) overlap of orbitals, stronger, allows free rotation. Pi: sideways (lateral) overlap of p-orbitals, weaker, restricts rotation. A double bond = 1 σ + 1 π; triple = 1 σ + 2 π.

List the three main types of intermolecular forces from weakest to strongest.

London dispersion forces (weakest) < dipole-dipole forces < hydrogen bonding (strongest of the three).

What conditions are required for hydrogen bonding?

An H atom covalently bonded to a highly electronegative atom (N, O, or F) interacting with a lone pair on N, O, or F of another molecule.

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

Water molecules form extensive hydrogen bonds (O-H...O), which require much more energy to break than the weaker dipole-dipole/dispersion forces in H2S, raising water's boiling point.

State Boyle's Law with its equation.

At constant temperature and amount of gas, volume is inversely proportional to pressure: P ∝ 1/V, or P1V1 = P2V2.

State Charles's Law with its equation.

At constant pressure and amount, volume is directly proportional to absolute temperature: V ∝ T, or V1/T1 = V2/T2 (T in kelvin).

State the ideal gas equation and define each term.

PV = nRT, where P = pressure, V = volume, n = moles, R = universal gas constant (8.314 J/mol·K or 0.0821 L·atm/mol·K), T = absolute temperature (K).

State Avogadro's Law and the molar volume of a gas at STP.

Avogadro's Law: equal volumes of gases at the same T and P contain equal numbers of molecules (V ∝ n). At STP one mole of any ideal gas occupies 22.4 L.

State Dalton's Law of Partial Pressures.

The total pressure of a mixture of non-reacting gases equals the sum of the partial pressures of the individual gases: P_total = P1 + P2 + P3 + ...

Compare the kinetic energy and molecular arrangement of solids, liquids, and gases.

Solids: low KE, fixed close-packed ordered arrangement. Liquids: moderate KE, close but mobile particles. Gases: high KE, far apart and random. KE increases solid < liquid < gas.

Distinguish crystalline from amorphous solids.

Crystalline solids have a regular, repeating long-range ordered arrangement and sharp melting points (e.g., NaCl). Amorphous solids lack long-range order and soften over a range (e.g., glass).

Define vapor pressure and the boiling point of a liquid.

Vapor pressure is the pressure exerted by a vapor in equilibrium with its liquid at a given temperature. Boiling point is the temperature at which vapor pressure equals the external (atmospheric) pressure.

Write the equilibrium constant expression Kc for aA + bB ⇌ cC + dD.

Kc = ([C]^c [D]^d) / ([A]^a [B]^b), using equilibrium molar concentrations. Pure solids and liquids are omitted.

What does the magnitude of K indicate about a reaction at equilibrium?

K >> 1: products favored (equilibrium lies to the right). K << 1: reactants favored (lies to the left). K ≈ 1: comparable amounts of reactants and products.

How are Kp and Kc related?

Kp = Kc(RT)^Δn, where Δn = (moles of gaseous products) − (moles of gaseous reactants), R = gas constant, T = temperature in kelvin.

State Le Chatelier's Principle.

If a system at equilibrium is subjected to a change in concentration, pressure, or temperature, the equilibrium shifts in the direction that counteracts (partially offsets) the imposed change.

According to Le Chatelier, how does increasing pressure affect a gaseous equilibrium?

Increasing pressure (decreasing volume) shifts the equilibrium toward the side with fewer moles of gas. If both sides have equal moles of gas, pressure has no effect.

How does increasing temperature shift exothermic vs endothermic equilibria?

Increasing temperature shifts an exothermic reaction toward reactants (K decreases) and an endothermic reaction toward products (K increases).

Does adding a catalyst shift the position of equilibrium?

No. A catalyst speeds up both forward and reverse reactions equally, so equilibrium is reached faster but the position and value of K are unchanged.

Define acids and bases according to the Brønsted-Lowry and Arrhenius theories.

Arrhenius: acid releases H+, base releases OH- in water. Brønsted-Lowry: acid is a proton (H+) donor, base is a proton acceptor.

Define pH and pOH, and state their relationship at 25°C.

pH = -log[H+]; pOH = -log[OH-]. At 25°C, pH + pOH = 14, and Kw = [H+][OH-] = 1×10^-14.

Calculate the pH of a 0.01 M HCl solution.

HCl is a strong acid, fully dissociated, so [H+] = 0.01 M = 1×10^-2 M. pH = -log(10^-2) = 2.

What is a buffer solution and what does it contain?

A buffer resists changes in pH on adding small amounts of acid or base. It contains a weak acid and its conjugate base (or a weak base and its conjugate acid).

Write the solubility product (Ksp) expression for a salt AxBy that dissolves as xA^+ + yB^-.

Ksp = [A^+]^x [B^-]^y, using the equilibrium concentrations of the ions in a saturated solution. Solids are not included.

How is the precipitation of a salt predicted using Ksp and the ionic product (Q)?

Compare Q (ionic product) with Ksp: if Q > Ksp a precipitate forms (supersaturated); Q = Ksp is saturated (equilibrium); Q < Ksp is unsaturated (no precipitate, more can dissolve).

What is the common ion effect?

The decrease in solubility of a slightly soluble salt when a soluble salt sharing a common ion is added; the added ion shifts the dissolution equilibrium backward (Le Chatelier).

Define enthalpy change (ΔH) and distinguish exothermic from endothermic reactions.

ΔH is the heat change at constant pressure. Exothermic: releases heat, ΔH negative. Endothermic: absorbs heat, ΔH positive.

State Hess's Law.

The total enthalpy change for a reaction is the same regardless of the route taken, depending only on initial and final states. So ΔH for an overall reaction equals the sum of ΔH of its steps.

Define standard enthalpy of formation (ΔHf°) and give its value for an element in its standard state.

The enthalpy change when one mole of a compound forms from its constituent elements in their standard states at 298 K and 1 atm. ΔHf° of any element in its standard state is zero.

How is ΔH°reaction calculated from standard enthalpies of formation?

ΔH°rxn = Σ ΔHf°(products) − Σ ΔHf°(reactants), each multiplied by its stoichiometric coefficient.

Define entropy (S) and state the sign of ΔS for melting and for gas formation.

Entropy is a measure of disorder/randomness in a system. ΔS is positive (increases) for melting (solid→liquid) and for forming gases or increasing gas moles.

Write the Gibbs free energy equation and the condition for spontaneity.

ΔG = ΔH − TΔS. A reaction is spontaneous when ΔG < 0, at equilibrium when ΔG = 0, and non-spontaneous when ΔG > 0.

Under what combinations of ΔH and ΔS signs is a reaction always or never spontaneous?

ΔH negative and ΔS positive: spontaneous at all temperatures. ΔH positive and ΔS negative: never spontaneous. The other two cases are temperature-dependent.

How is rate of reaction defined and what are its typical units?

Rate is the change in concentration of a reactant or product per unit time (Δ[concentration]/Δtime). Units: mol dm^-3 s^-1 (or M/s).

What is a rate law, and define the order of reaction.

Rate law: Rate = k[A]^m[B]^n, where k is the rate constant. The order with respect to a reactant is its exponent; overall order = sum of exponents (m + n).

What is the half-life of a first-order reaction and how does it depend on concentration?

For first order, t1/2 = 0.693/k. It is constant and independent of the initial concentration of reactant.

List the main factors that affect the rate of a chemical reaction.

Concentration (or pressure for gases), temperature, surface area of solids, presence of a catalyst, and the nature of reactants (and light for some reactions).

According to collision theory, what two conditions must colliding molecules satisfy to react?

They must collide with sufficient energy (≥ activation energy, Ea) and with the correct (proper) orientation.

Why does increasing temperature increase reaction rate?

Higher temperature increases the average kinetic energy, so a greater fraction of molecules have energy ≥ Ea and collisions are more frequent and more energetic, increasing the rate (roughly doubling per 10°C rise).

Define activation energy (Ea).

The minimum energy that colliding reactant molecules must possess to form the activated complex and convert into products.

What is a catalyst and how does it speed up a reaction?

A substance that increases reaction rate without being consumed, by providing an alternative reaction pathway with a lower activation energy. It does not change ΔH or the equilibrium position.

Differentiate homogeneous and heterogeneous catalysis with an example of each.

Homogeneous: catalyst in the same phase as reactants (e.g., H2SO4 in ester hydrolysis). Heterogeneous: catalyst in a different phase (e.g., solid iron in the Haber process with gaseous reactants).