MDCAT Chemistry Syllabus

State Graham's law of diffusion.

The rate of diffusion of a gas is inversely proportional to the square root of its density (or molar mass): $$\frac{r_1}{r_2} = \sqrt{\frac{M_2}{M_1}}$$

List the main properties of liquids compared to gases and solids.

Liquids have a definite volume but no definite shape, are nearly incompressible, have intermolecular forces intermediate between gases and solids, and possess fluidity, surface tension, and vapour pressure.

Define vapour pressure and evaporation.

Vapour pressure is the pressure exerted by the vapour in equilibrium with its liquid at a given temperature. Evaporation is the escape of surface molecules from a liquid into the vapour phase below the boiling point.

How does boiling point relate to vapour pressure and external pressure?

A liquid boils when its vapour pressure equals the external (atmospheric) pressure. Higher external pressure raises the boiling point; lower external pressure lowers it.

Distinguish crystalline and amorphous solids.

Crystalline solids have a regular, long-range ordered arrangement, sharp melting points, and anisotropic properties. Amorphous solids have irregular short-range order, melt over a range, and are isotropic (e.g. glass).

What is an isomorphous and a polymorphous substance?

Isomorphism: two substances with different chemical compositions but the same crystalline form. Polymorphism: one substance existing in more than one crystalline form (e.g. carbon as diamond and graphite).

State the First Law of Thermodynamics and its equation.

Energy can neither be created nor destroyed, only converted. $$\Delta E = q + w$$ where $\Delta E$ is change in internal energy, $q$ is heat added to the system, and $w$ is work done on the system.

Distinguish between a state function and a path function, with examples.

A state function depends only on initial and final states (e.g. $\Delta E$, $\Delta H$, entropy). A path function depends on the route taken (e.g. heat $q$ and work $w$).

Define enthalpy and relate it to internal energy.

Enthalpy is the heat content at constant pressure: $H = E + PV$. The change is $\Delta H = \Delta E + P\Delta V$; at constant pressure $\Delta H = q_p$.

State the Second Law of Thermodynamics in terms of entropy.

The entropy of the universe always increases for a spontaneous (irreversible) process: $$\Delta S_{\text{universe}} > 0$$

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

$$\Delta G = \Delta H - T\Delta S$$ A process is spontaneous when $\Delta G < 0$, at equilibrium when $\Delta G = 0$, and non-spontaneous when $\Delta G > 0$.

Distinguish exothermic and endothermic reactions by the sign of $\Delta H$.

Exothermic reactions release heat, $\Delta H < 0$ (negative). Endothermic reactions absorb heat, $\Delta H > 0$ (positive).

Define enthalpy (heat) of formation.

The enthalpy change when one mole of a compound is formed from its constituent elements in their standard states, denoted $\Delta H_f$.

State Hess's law of constant heat summation.

The total enthalpy change of a reaction is independent of the pathway taken, depending only on the initial and final states. Enthalpies of individual steps can be added to give the overall $\Delta H$.

Define enthalpy of neutralisation and give its typical value for a strong acid and strong base.

The heat released when one mole of water is formed by the neutralisation of an acid by a base. For a strong acid and strong base it is approximately $-57.3\ \text{kJ mol}^{-1}$.

What is a reversible reaction and what characterises dynamic equilibrium?

A reversible reaction proceeds in both forward and backward directions. At dynamic equilibrium the forward and reverse reaction rates are equal, so the concentrations of reactants and products remain constant.

Write the equilibrium constant expression $K_c$ for $\ce{aA + bB <=> cC + dD}$.

$$K_c = \frac{[\ce{C}]^{c}[\ce{D}]^{d}}{[\ce{A}]^{a}[\ce{B}]^{b}}$$

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) that change.

Using Le Chatelier's principle, predict the effect of increasing pressure on the equilibrium $\ce{N2 + 3H2 <=> 2NH3}$.

Increasing pressure shifts the equilibrium toward the side with fewer moles of gas, i.e. toward $\ce{NH3}$ (4 moles $\to$ 2 moles), increasing ammonia yield.

How does a catalyst affect a reaction at equilibrium?

A catalyst speeds up forward and reverse reactions equally, so it reduces the time to reach equilibrium but does NOT change the equilibrium position or the value of $K_c$.

State the Bronsted-Lowry definitions of acid and base.

An acid is a proton ($\ce{H+}$) donor and a base is a proton acceptor.

Define pH and pOH and give their relationship at $25\,^{\circ}\text{C}$.

$\text{pH} = -\log[\ce{H+}]$ and $\text{pOH} = -\log[\ce{OH-}]$. At $25\,^{\circ}\text{C}$: $$\text{pH} + \text{pOH} = 14$$

What is the ionic product of water $K_w$ and its value at $25\,^{\circ}\text{C}$?

$$K_w = [\ce{H+}][\ce{OH-}] = 1\times10^{-14}\ \text{mol}^{2}\,\text{dm}^{-6}$$

Define a Lewis acid and a Lewis base.

A Lewis acid is an electron-pair acceptor; a Lewis base is an electron-pair donor.

What is a buffer solution and give an example of an acidic buffer?

A buffer resists changes in pH on adding small amounts of acid or base. An acidic buffer is a mixture of a weak acid and its salt, e.g. $\ce{CH3COOH}$ and $\ce{CH3COONa}$.

What is catenation and why is carbon exceptional at it?

Catenation is the self-linking of atoms of the same element into chains and rings. Carbon excels due to its strong, stable $\ce{C-C}$ bonds and tetravalency, allowing vast numbers of organic compounds.

Define structural isomerism and name its types.

Structural isomers have the same molecular formula but different structural arrangements. Types: chain, position, functional group, and metamerism isomerism.

Distinguish alkanes, alkenes, and alkynes by their general formulas and bonding.

Alkanes: $C_nH_{2n+2}$, single $\ce{C-C}$ bonds (saturated). Alkenes: $C_nH_{2n}$, one $\ce{C=C}$ double bond. Alkynes: $C_nH_{2n-2}$, one $\ce{C#C}$ triple bond.

Compare the characteristic reactions of alkanes and alkenes.

Alkanes are saturated and undergo substitution reactions (e.g. free-radical halogenation). Alkenes are unsaturated and undergo addition reactions (e.g. with $\ce{H2}$, $\ce{Br2}$, $\ce{HX}$).

State Markownikoff's rule for addition of $\ce{HX}$ to an unsymmetrical alkene.

When $\ce{HX}$ adds to an unsymmetrical alkene, the hydrogen atom attaches to the doubly bonded carbon already bearing more hydrogen atoms, and the halogen to the carbon with fewer hydrogens.

Identify the functional groups: $\ce{-OH}$, $\ce{-CHO}$, $\ce{-COOH}$, and $\ce{>C=O}$.

$\ce{-OH}$: hydroxyl (alcohols). $\ce{-CHO}$: aldehyde group. $\ce{-COOH}$: carboxyl (carboxylic acids). $\ce{>C=O}$ (between two carbons): carbonyl (ketones).