HKDSE Chemistry flashcards

32 cards. Read the term and the definition here. The deck at the top of this page is optional.

Mole definition

One mole contains 6.022×10^23 particles (Avogadro constant). It is the amount of substance whose mass in grams equals its relative atomic or molecular mass.

Molar volume of gas

At room temperature and pressure, one mole of any gas occupies about 24.0 dm^3; at s.t.p. it occupies about 22.4 dm^3. Used to convert gas volume to moles.

Avogadro's law

Equal volumes of different gases at the same temperature and pressure contain equal numbers of moles. Hence gas volume ratios in reactions equal mole ratios.

Empirical formula

The simplest whole-number ratio of atoms of each element in a compound. Found by dividing each element's mass by its relative atomic mass and simplifying.

Molecular formula

The actual number of atoms of each element in one molecule. It equals the empirical formula multiplied by a whole-number factor obtained from the molar mass.

Limiting reactant

The reactant completely consumed first, which stops the reaction. It is identified by calculating the moles of product each reactant could form and taking the smaller.

Percentage yield

Actual amount of product obtained divided by theoretical maximum amount, expressed as a percentage. Values below 100% arise from side reactions, incomplete reaction and losses.

Ideal gas equation

PV = nRT, where P is pressure in Pa, V is volume in m^3, n is moles, T is temperature in kelvin and R is 8.314 J K^-1 mol^-1.

Concentration formula

Concentration in mol dm^-3 equals moles of solute divided by volume of solution in dm^3. On dilution, c1V1 = c2V2 because solute moles stay constant.

Oxidation number rules

Oxidation numbers sum to the overall charge. Oxygen is usually −2, hydrogen +1, group 1 metals +1, group 2 metals +2, and free elements are 0.

Oxidation and reduction

Oxidation is an increase in oxidation number (loss of electrons); reduction is a decrease (gain of electrons). A redox reaction involves both simultaneously.

Ionic bonding

Electrostatic attraction between oppositely charged ions in a giant lattice, typically formed between metals and non-metals. High melting point; conducts when molten or aqueous.

Giant covalent structure

A three-dimensional network of atoms joined by covalent bonds, as in diamond, graphite and silicon dioxide. Very high melting point; usually insoluble; generally non-conducting except graphite.

Simple molecular substance

Discrete molecules held together by weak intermolecular forces. Low melting and boiling points; does not conduct electricity because there are no free ions or electrons.

Metallic bonding

A lattice of positive metal ions surrounded by a sea of delocalised electrons. Explains electrical and thermal conductivity, malleability and ductility of metals.

First ionisation energy

The energy needed to remove one mole of electrons from one mole of gaseous atoms: X(g) → X^+(g) + e^-. It increases across a period and decreases down a group.

Electronegativity

The tendency of an atom to attract electrons in a covalent bond. It increases across a period and decreases down a group; large differences give polar bonds.

Collision theory

Particles must collide with energy at least equal to the activation energy and with correct orientation for reaction. Raising temperature increases both collision frequency and energy.

Catalyst

A substance that increases reaction rate by providing an alternative pathway with lower activation energy. It is not consumed and does not change the equilibrium position.

Le Chatelier's principle

When a system at equilibrium is disturbed by a change in concentration, pressure or temperature, it shifts in the direction that opposes the change.

Equilibrium constant Kc

For aA + bB ⇌ cC + dD, Kc = [C]^c[D]^d / ([A]^a[B]^b). It depends only on temperature; a catalyst does not alter it.

pH definition

pH = −log[H^+]. At 25 °C, Kw = [H^+][OH^-] = 1.0×10^-14, so pH + pOH = 14 for any aqueous solution.

Strong acid

An acid that ionises completely in water, such as HCl, HNO3 and H2SO4. Its pH depends on the acid concentration and the number of ionisable hydrogen atoms.

Titration calculation

Moles = concentration × volume in dm^3. Convert moles of titrant to moles of analyte using the balanced equation ratio, then divide by the analyte volume.

Standard enthalpy change

The enthalpy change when one mole of substance is formed from elements in standard states, or when one mole of fuel is completely burned, at 298 K and 1 atm.

Hess's law

The total enthalpy change of a reaction is independent of the route taken. It allows ΔH to be calculated from formation or combustion enthalpy cycles.

Bond breaking and forming

Breaking bonds absorbs energy (endothermic); forming bonds releases energy (exothermic). ΔH ≈ energy needed to break bonds − energy released forming bonds.

Electrolysis

Decomposition of an ionic compound by passing direct current through it molten or in aqueous solution. Cations move to the cathode and are reduced; anions move to the anode.

Test for carbon dioxide

Bubble the gas through limewater (calcium hydroxide solution). Carbon dioxide turns it milky or cloudy because insoluble calcium carbonate is formed.

Test for alkenes

Shake with bromine dissolved in an organic solvent or with bromine water. Alkenes decolourise it rapidly by addition; alkanes do not react in the dark.

Esterification

A carboxylic acid reacts with an alcohol in the presence of concentrated sulfuric acid catalyst to give an ester and water. The reaction is reversible.

Homologous series

A family of organic compounds with the same general formula and functional group, showing similar chemical properties and a gradual trend in physical properties.