IGCSE Chemistry

Solid

A state of matter with a fixed shape and fixed volume, where constituent particles vibrate about fixed positions within a regular lattice.

Liquid

A state of matter with a fixed volume that conforms to the shape of its container, in which particles remain closely packed but are free to flow past one another.

Gas

A state of matter possessing neither fixed shape nor fixed volume, characterized by particles that are spaced widely apart and moving rapidly in random directions.

Melting

The phase transition from a solid to a liquid occurring at a characteristic melting point upon absorbing thermal energy.

Freezing

The phase transition from a liquid to a solid occurring when a substance releases thermal energy.

Boiling

The phase transition from a liquid to a gas occurring rapidly throughout the entire liquid body at a fixed boiling point.

Evaporation

The phase transition from a liquid to a gas taking place exclusively at the liquid surface at temperatures below the boiling point.

Condensation

The phase transition from a gas to a liquid resulting from the removal of thermal energy.

Sublimation

The direct phase change of a substance from a solid into a gas without passing through an intermediate liquid state.

Deposition

The direct phase change of a substance from a gas into a solid without entering a liquid phase.

Diffusion

The net movement of particles from an area of higher concentration to an area of lower concentration down a concentration gradient.

Brownian motion

The random, erratic motion of microscopic particles suspended in a fluid resulting from continuous collisions with surrounding fluid molecules.

Absolute zero

The theoretical temperature at which particles possess zero kinetic energy, defined as $-273\,^{\circ}\text{C}$ or $0\,\text{K}$.

Effect of molecular mass on diffusion

Gases with lower relative molecular masses ($M_r$) diffuse at a faster rate than heavier gases at the same temperature.

Atom

The smallest electrically neutral unit of an element that can take part in a chemical reaction.

Element

A pure substance composed of only one type of atom that cannot be broken down into simpler substances by chemical means.

Compound

A pure substance composed of two or more different elements chemically bonded together in fixed stoichiometric ratios.

Mixture

A combination of two or more distinct substances physically intermingled without being chemically bonded.

Proton

A subatomic particle located in the nucleus with a relative mass of $1$ and a relative electrical charge of $+1$.

Neutron

A neutral subatomic particle situated in the nucleus with a relative mass of $1$ and a relative electrical charge of $0$.

Electron

A subatomic particle orbiting the nucleus in energy levels, possessing a negligible relative mass of $\frac{1}{1840}$ and a relative charge of $-1$.

Atomic number

The number of protons contained within the nucleus of an atom, establishing its elemental identity.

Mass number

The total sum of protons and neutrons located inside the atomic nucleus.

Isotopes

Atoms belonging to the exact same element that contain identical numbers of protons but differing numbers of neutrons.

Relative atomic mass ($A_r$)

The weighted average mass of naturally occurring isotopes of an element compared to $\frac{1}{12}$ of the mass of a carbon-12 atom.

Relative molecular mass ($M_r$)

The sum of the relative atomic masses ($A_r$) of all atoms present in a molecular formula.

Valence electrons

Electrons residing in the outermost energy level or shell of an atom.

Cation

A positively charged ion formed when a neutral atom loses one or more electrons.

Anion

A negatively charged ion formed when a neutral atom gains one or more electrons.

Ionic bonding

The strong electrostatic attraction operating between oppositely charged ions formed by the transfer of electrons.

Giant ionic lattice

A regular, three-dimensional repeating array of alternating positive and negative ions held firmly by electrostatic attractions.

Electrical conductivity of ionic compounds

Non-conductive in the solid state because ions are locked in position, but conductive when molten or aqueous because ions become free to migrate.

Covalent bonding

Electrostatic attraction between a shared pair of electrons and the positively charged nuclei of the bonded non-metal atoms.

Simple molecular structure

Discrete groups of covalently bonded atoms held to neighbouring molecules by relatively weak intermolecular forces.

Giant covalent structure

A vast network of non-metal atoms joined together entirely by strong covalent bonds extending through the whole crystal.

Diamond

An allotrope of carbon in which each carbon atom is covalently bonded to four others in a rigid tetrahedral lattice, conferring high hardness.

Graphite

An allotrope of carbon arranged in hexagonal parallel layers with delocalized electrons that confer electrical conductivity.

Silicon dioxide ($SiO_2$)

A giant covalent network in which each silicon atom bonds covalently to four oxygen atoms and each oxygen atom to two silicon atoms.

Metallic bonding

The electrostatic attraction between a lattice of positive metal ions and a surrounding sea of delocalized electrons.

Malleability and ductility of metals

Mechanical flexibility resulting from regular layers of metal cations sliding over one another without disrupting the metallic bond.

Alloy

A mixture of a base metal with one or more other elements, engineered to enhance strength, hardness, or corrosion resistance.

Structural cause of alloy hardness

Atoms of differing atomic radii disrupt the uniform crystal lattice layers, impeding them from sliding smoothly past one another.

The mole

The amount of a chemical substance containing precisely $6.02 \times 10^{23}$ specified constituent entities.

Avogadro constant

The fundamental constant representing the number of particles in one mole of any substance, defined as $6.02 \times 10^{23}\,mol^{-1}$.

Relationship between moles, mass, and molar mass

$$\text{Moles} = \frac{\text{Mass}\,(g)}{M_r}$$

Formula for solution concentration

$$\text{Concentration}\,(mol\,dm^{-3}) = \frac{\text{Moles}}{\text{Volume}\,(dm^3)}$$

Molar gas volume at room temperature and pressure

$$24\,dm^3\,mol^{-1}$$

Empirical formula

The simplest whole-number ratio of the atoms of each element present in a chemical compound.

Molecular formula

The actual number of atoms of each individual element present within a single molecule of a compound.

Percentage yield formula

$$\text{Percentage Yield} = \left(\frac{\text{Actual Yield}}{\text{Theoretical Yield}}\right) \times 100$$

Limiting reactant

The reactant in a chemical process that is entirely consumed first, thereby capping the theoretical yield of products.

Exothermic reaction

A chemical reaction that transfers thermal energy to the surroundings, leading to a rise in ambient temperature and a negative $\Delta H$.

Endothermic reaction

A chemical reaction that absorbs thermal energy from the surroundings, resulting in a temperature drop and a positive $\Delta H$.

Activation energy ($E_a$)

The minimum kinetic energy that colliding reactant particles must possess for a collision to result in a chemical reaction.

Calculating enthalpy change from bond energies

$$\Delta H = \Sigma(\text{Bond Energies Broken}) - \Sigma(\text{Bond Energies Formed})$$

Energetics of bond breaking versus bond making

Bond breaking is an endothermic process requiring energy input, whereas bond formation is an exothermic process releasing energy.

Electrolysis

The chemical decomposition of an ionic compound, either molten or dissolved in water, driven by the passage of a direct electric current.

Electrolyte

A molten ionic compound or aqueous solution containing mobile ions that conducts electricity during electrolysis.

Cathode

The negative electrode of an electrolytic cell at which positively charged cations undergo reduction by gaining electrons.

Anode

The positive electrode of an electrolytic cell at which negatively charged anions undergo oxidation by losing electrons.

Products of molten lead(II) bromide electrolysis

Lead metal produced at the cathode ($Pb^{2+} + 2e^- \rightarrow Pb$) and red-brown bromine gas produced at the anode ($2Br^- \rightarrow Br_2 + 2e^-$).

Products of concentrated aqueous sodium chloride electrolysis

Hydrogen gas at the cathode, chlorine gas at the anode, and sodium hydroxide left in the aqueous electrolyte.

Discharge rule at the cathode in aqueous electrolysis

Hydrogen gas is liberated unless the metal cation is less reactive than hydrogen, in which case the metal itself plates out.

Discharge rule at the anode in aqueous electrolysis

Halogen gas evolves if halide ions are sufficiently concentrated; otherwise, oxygen gas is generated via hydroxide ion oxidation.

Electroplating setup requirements

The object to be plated acts as the cathode, the coating metal serves as the anode, and the electrolyte contains cations of the coating metal.

Collision theory

The principle stating that reacting particles must collide with energy greater than or equal to the activation energy and with appropriate spatial orientation.

Mechanism of a catalyst

Provides an alternative reaction route characterized by a lower activation energy ($E_a$), increasing the proportion of successful collisions.

Dynamic equilibrium

A state achieved in a closed system where the rate of the forward reaction equals the rate of the backward reaction, keeping concentrations constant.

Le Chatelier's principle

If a system existing at dynamic equilibrium is subjected to an external perturbation, the equilibrium position shifts to oppose that change.

Temperature change on dynamic equilibrium

An increase in temperature displaces equilibrium toward the endothermic direction, whereas a temperature drop favors the exothermic direction.

Pressure change on gas equilibrium

An increase in pressure shifts equilibrium toward the side having fewer gas molecules; a decrease shifts toward more gas molecules.

Haber process operating conditions

A temperature of $450\,^{\circ}\text{C}$, a pressure of $200\,\text{atm}$, and an iron catalyst for the synthesis of ammonia ($N_2 + 3H_2 \rightleftharpoons 2NH_3$).

Contact process operating conditions

A temperature of $450\,^{\circ}\text{C}$, a pressure of $1\text{ to }2\,\text{atm}$, and a vanadium(V) oxide catalyst ($V_2O_5$) for $2SO_2 + O_2 \rightleftharpoons 2SO_3$.

Alkali metals reactivity trend

Chemical reactivity increases down Group 1 because valence electrons sit further from nuclear pull and are lost more readily.

Halogens reactivity trend

Chemical reactivity decreases down Group 7 as increased atomic radius and electron shielding weaken the attraction for an incoming electron.

Physical states and colors of halogens at room temperature

Fluorine is a pale yellow gas, chlorine is a pale green gas, bromine is a red-brown liquid, and iodine is a dark grey solid.

Noble gases inertness

High chemical stability due to possession of a complete valence electron shell, removing any thermodynamic drive to react.

Transition elements characteristic properties

High densities, high melting points, variable oxidation states, production of colored ionic complexes, and significant catalytic activity.

Rusting essential conditions

Iron and steel corrosion requires the simultaneous presence of both water and oxygen.

Sacrificial protection against corrosion

Affixing a more reactive metal, such as zinc or magnesium, which oxidizes preferentially in place of iron.

Hydrocarbon

An organic chemical compound containing exclusively carbon and hydrogen atoms.

Alkanes general formula and saturation

Saturated hydrocarbons containing solely single carbon-carbon bonds, adhering to the general formula $C_n H_{2n+2}$.

Alkenes general formula and functional group

Unsaturated hydrocarbons featuring a carbon-carbon double bond ($C=C$), characterized by the general formula $C_n H_{2n}$.

Chemical test for unsaturation

Aqueous bromine water transitions from an orange-brown color to completely colorless in the presence of carbon-carbon double bonds.

Cracking

The thermal breakdown of long-chain alkanes into smaller, commercially useful alkanes and unsaturated alkenes using heat and a catalyst.

Crude oil fractions sequence by boiling point

Refinery gas, gasoline (petrol), kerosene (paraffin), diesel oil, fuel oil, and bitumen residue.

Test for hydrogen gas

Exposing the gas to a lighted splint produces an audible squeaky pop sound.

Test for oxygen gas

Introducing a glowing wooden splint into the gas causes the splint to relight.

Test for carbon dioxide gas

Bubbling the gas through an aqueous solution of calcium hydroxide (limewater) causes it to turn milky or cloudy.

Test for chlorine gas

Exposing the gas to damp blue litmus paper causes the paper to turn temporarily red before being bleached white.

Flame test colors for common cations

Lithium gives a crimson-red flame, sodium gives a persistent yellow flame, potassium gives a lilac flame, and copper(II) gives a blue-green flame.

Test for sulfate ions ($SO_4^{2-}$)

Acidifying with dilute hydrochloric acid followed by adding aqueous barium chloride produces an insoluble white precipitate of barium sulfate.

Test for halide ions ($Cl^-$, $Br^-$, $I^-$)

Acidifying with dilute nitric acid followed by aqueous silver nitrate produces white ($AgCl$), cream ($AgBr$), or yellow ($AgI$) precipitates.