Air University Entry Test Chemistry Syllabus

How does electronegativity difference classify a bond as ionic, polar covalent, or non-polar covalent?

Difference > ~1.7 is largely ionic; ~0.4 to 1.7 is polar covalent; < ~0.4 is non-polar covalent.

What is a coordinate (dative) covalent bond?

A covalent bond in which both shared electrons are donated by the same atom, e.g., the N-to-H bond in the ammonium ion NH4+ formed from NH3 and H+.

State the octet rule.

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

What is hybridization?

The mixing of atomic orbitals (e.g., s and p) on an atom to form an equal number of new, equivalent hybrid orbitals suited for bonding and a particular geometry.

Give the geometry and bond angle for sp, sp2, and sp3 hybridization.

sp = linear, 180 degrees; sp2 = trigonal planar, 120 degrees; sp3 = tetrahedral, 109.5 degrees.

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

sp3 hybridized, tetrahedral geometry with bond angles of 109.5 degrees.

According to VSEPR theory, what shapes result from 2, 3, and 4 bonding electron pairs with no lone pairs?

2 pairs = linear, 3 pairs = trigonal planar, 4 pairs = tetrahedral. Electron pairs arrange to minimize repulsion.

Why is the bond angle in water (104.5 degrees) less than the tetrahedral angle (109.5 degrees)?

Oxygen has two lone pairs whose lone-pair/lone-pair and lone-pair/bond-pair repulsions are stronger than bond-pair/bond-pair repulsion, compressing the H-O-H angle below 109.5 degrees.

Distinguish between sigma and pi bonds.

A sigma bond forms by head-on (axial) overlap of orbitals and is stronger; a pi bond forms by sideways (lateral) overlap of p-orbitals and is weaker. A double bond = 1 sigma + 1 pi; a triple bond = 1 sigma + 2 pi.

List the three main types of intermolecular forces in order of increasing strength.

London dispersion (induced dipole) forces < dipole-dipole forces < hydrogen bonding (the strongest).

What conditions are required for hydrogen bonding?

A hydrogen atom must be bonded to a highly electronegative atom (N, O, or F), and there must be a lone pair on an N, O, or F atom in a neighboring molecule.

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

Water molecules form extensive hydrogen bonds (O is far more electronegative than S), requiring much more energy to separate, whereas H2S has only weak dipole-dipole and dispersion forces.

What is the effect of molecular size/molar mass on London dispersion forces?

Dispersion forces increase with increasing number of electrons (larger molar mass / larger molecules) because larger electron clouds are more polarizable, raising boiling points.

State Boyle's law.

At constant temperature and amount of gas, the volume of a gas is inversely proportional to its pressure: PV = constant, or P1V1 = P2V2.

State Charles's law.

At constant pressure and amount of gas, the volume of a gas is directly proportional to its absolute (Kelvin) temperature: V/T = constant, or V1/T1 = V2/T2.

Write the ideal gas equation and define each term.

PV = nRT, where P = pressure, V = volume, n = moles, R = universal gas constant (0.0821 L atm / mol K, or 8.314 J/mol K), 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; volume is directly proportional to the number of moles (V/n = constant).

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 + ...

State Graham's law of diffusion.

The rate of diffusion (or effusion) of a gas is inversely proportional to the square root of its molar mass: rate proportional to 1/sqrt(M). Lighter gases diffuse faster.

Under what conditions do real gases deviate most from ideal behavior?

At high pressure and low temperature, where molecular volume and intermolecular attractions become significant and the ideal gas assumptions break down.

Define vapor pressure of a liquid.

The pressure exerted by a vapor in equilibrium with its liquid in a closed container at a given temperature. It increases as temperature increases.

Define the boiling point of a liquid.

The temperature at which a liquid's vapor pressure equals the external (atmospheric) pressure, so vaporization occurs throughout the liquid.

What is surface tension, and why does it occur?

Surface tension is the energy/force that minimizes a liquid's surface area, arising because surface molecules experience a net inward attraction (unbalanced intermolecular forces). Stronger intermolecular forces give higher surface tension.

How does temperature affect the viscosity of a liquid?

Viscosity (resistance to flow) decreases as temperature increases, because molecules gain kinetic energy and overcome intermolecular attractions more easily.

What are the seven crystal systems based on?

They are classified by the relative lengths of the unit-cell axes (a, b, c) and the angles between them: cubic, tetragonal, orthorhombic, monoclinic, triclinic, hexagonal, and rhombohedral (trigonal).

Compare crystalline and amorphous solids.

Crystalline solids have a regular, long-range ordered arrangement of particles, sharp melting points, and definite geometry. Amorphous solids (e.g., glass) lack long-range order and soften over a range of temperatures.

How many atoms per unit cell are in simple cubic, body-centered cubic (BCC), and face-centered cubic (FCC) lattices?

Simple cubic = 1 atom; BCC = 2 atoms; FCC = 4 atoms per unit cell.

Name the four types of crystalline solids and give a property of each.

Ionic (high melting, brittle, e.g., NaCl); covalent network (very hard, very high melting, e.g., diamond); metallic (conduct electricity, malleable); molecular (soft, low melting, e.g., ice).

State the first law of thermodynamics.

Energy can neither be created nor destroyed, only converted from one form to another. ΔU = q + w (change in internal energy = heat added + work done on the system).

Define enthalpy of reaction (ΔH) and distinguish exothermic from endothermic.

ΔH is the heat change at constant pressure. Exothermic reactions release heat (ΔH negative); endothermic reactions absorb heat (ΔH positive).

State Hess's law.

The total enthalpy change of a reaction is the same regardless of the route taken, depending only on the initial and final states. Enthalpy changes of steps can be added to give the overall ΔH.

What is the standard enthalpy of formation (ΔHf)?

The enthalpy change when one mole of a compound is formed from its elements in their standard states. The ΔHf of any element in its standard state is zero.

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 (opposes) the change.

Write the equilibrium constant expression Kc for the reaction aA + bB <=> cC + dD.

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

What does the magnitude of the equilibrium constant indicate?

A large K (>>1) means products are favored at equilibrium; a small K (<<1) means reactants are favored; K ~ 1 means comparable amounts of both.

For the exothermic reaction N2 + 3H2 <=> 2NH3, how do increased pressure and decreased temperature affect the yield of ammonia?

Increased pressure shifts equilibrium toward fewer gas moles (toward NH3), increasing yield. Decreased temperature favors the exothermic forward reaction, also increasing NH3 yield.

Does a catalyst change the position of equilibrium or the value of K?

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

Define the rate of a chemical reaction.

The change in concentration of a reactant or product per unit time, e.g., mol L^-1 s^-1 (decrease for reactants, increase for products).

List the main factors that affect reaction rate.

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

How does a catalyst increase reaction rate?

It provides an alternative reaction pathway with a lower activation energy, allowing more colliding molecules to react, without being consumed in the reaction.

State the collision theory of reaction rates.

Reactions occur when reactant particles collide with sufficient energy (equal to or greater than the activation energy) and with the correct orientation. Increasing collision frequency or energy increases the rate.

What is activation energy?

The minimum energy that colliding reactant particles must possess for a reaction to occur (to form the activated complex/transition state).

Why does increasing temperature increase reaction rate?

It increases the average kinetic energy of particles, so collisions are more frequent and a larger fraction of molecules have energy equal to or above the activation energy.

What is the order of reaction?

The sum of the powers of the concentration terms in the experimentally determined rate law (rate = k[A]^x[B]^y; order = x + y). It indicates how rate depends on reactant concentrations.