A-Level Chemistry flashcards

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240 free AQA A-Level Chemistry flashcards across 3 topics. Every card is a key term and its definition, taken word for word from a published Revise Right lesson, and links back to that lesson so you can read the full notes behind it.

Physical chemistry

102 cards from 65 lessons. Each term links to the lesson it comes from.

Unveiling the Atom: A Journey Through Time

• Nuclear model of the atom: A model in which the atom consists of a tiny, dense, positively charged nucleus at the centre, surrounded by electrons in the largely empty space around it. The nucleus contains almost all the mass of the atom.
• Atomic number (proton number): The number of protons in the nucleus of an atom. It uniquely identifies the element and, in a neutral atom, equals the number of electrons. Symbol: Z.

Atomic Structure: Mass Number, Isotopes, and Calculations

• Atomic number (Z): The number of protons in the nucleus of an atom. Because every proton carries a charge of +1, Z also defines the nuclear charge. No two elements share the same atomic number — it is the identity card of an element.
• Mass number (A): The total number of protons and neutrons in the nucleus. Because electrons have negligible mass compared with nucleons, the mass number gives a good whole-number approximation of the atomic mass.

Mass Spectrometry: Principles and Interpretation

• Ionisation (TOF): In electrospray ionisation, the sample is dissolved and sprayed through a charged needle. Each molecule gains a proton (H+) to form a positive ion, typically [M+H]+. In electron ionisation, high-energy electrons knock one electron from each atom or molecule, forming a positive ion with a 1+ charge.
• Relative atomic mass (Ar): The weighted mean mass of an atom of an element, relative to one-twelfth of the mass of a carbon-12 atom. Calculated from a mass spectrum using: Ar = (sum of [isotopic mass × percentage abundance]) ÷ 100.

Ionisation Energy Explained

• First ionisation energy (IE1): The energy required to remove one mole of electrons from one mole of gaseous atoms in their ground state, forming one mole of gaseous 1+ ions. Measured in kJ mol⁻¹.

Electron Configuration Evidence

• First ionisation energy: The energy required to remove one mole of electrons from one mole of gaseous atoms in their ground state, producing one mole of gaseous 1+ ions. Equation: X(g) → X⁺(g) + e⁻. Measured in kJ mol⁻¹.

Balancing Equations and Stoichiometric Calculations

• Full equation: A full equation shows all reactants and products with correct formulae, balancing numbers, and state symbols. Unlike an ionic equation, it keeps soluble ionic compounds written as full formulae rather than splitting them into separate ions.

Percentage Yield and Atom Economy (Part 1/2)

• Percentage yield: Percentage yield = (actual yield / theoretical yield) × 100. It compares the mass of product actually collected to the maximum mass that stoichiometry predicts, expressed as a percentage.

Percentage Yield and Atom Economy (Part 2/2)

• Atom Economy: The percentage of the total mass of all reactants that is converted into the desired product. Formula: (molecular mass of desired product / sum of molecular masses of all reactants) × 100.

Relative Atomic Mass and Relative Molecular Mass

• Relative atomic mass (Ar): The weighted mean mass of an atom of an element, measured on a scale on which the mass of a carbon-12 atom is exactly 12. It is a dimensionless quantity — it has no units.
• Relative molecular mass (Mr): The weighted mean mass of a molecule relative to one-twelfth of the mass of a carbon-12 atom. Calculated by adding together the Ar values of all atoms present in the molecular formula. It is dimensionless.

Introduction to the Ideal Gas Equation

• Ideal gas: A theoretical gas whose particles have negligible volume and experience no intermolecular forces. Real gases approximate ideal behaviour at low pressures and high temperatures.

The Mole & Avogadro's Constant

• The Mole: The mole is the SI unit of amount of substance. One mole of a substance contains 6.022 × 10^23 elementary particles of that substance. The symbol for the mole is mol.
• Avogadro's Constant (L): Avogadro's constant is the number of particles in exactly one mole of any substance. Its value is 6.022 × 10^23 mol^-1. It applies equally to atoms, molecules, ions, electrons, and formula units.

Moles, Mass & Mr Calculations

• Relative Molecular Mass (Mr): The sum of the relative atomic masses (Ar) of all atoms in one molecule of a substance. It is dimensionless — it has no units. For ionic compounds, the same calculation applies but the quantity is called the relative formula mass, also given the symbol Mr.

Concentration Calculations

• Concentration (c): The amount of solute (in mol) dissolved in 1 dm³ of solution. Concentration is measured in mol dm⁻³ at A Level. A solution with a concentration of 2 mol dm⁻³ contains 2 moles of solute in every 1 dm³ of solution.

Predicting Ion Charges & Formula Construction

• Cation: A positively charged ion formed when an atom loses one or more electrons. The charge equals the number of electrons lost. Formed by metals and hydrogen.
• Anion: A negatively charged ion formed when an atom gains one or more electrons. The charge equals the number of electrons gained. Formed by non-metals.

Ionic Lattices & Compound Ions

• Ionic bonding: The strong electrostatic attraction between oppositely charged ions in a lattice. It arises after electron transfer from a metal atom to a non-metal atom, producing a positively charged cation and a negatively charged anion.

Crystal Structures: Ionic, Metallic & Molecular (Part 1/2)

• Ionic lattice: A giant, regular, three-dimensional arrangement of oppositely charged ions held together by strong electrostatic attractions acting in all directions throughout the structure. There are no discrete molecules — the formula unit (e.g. NaCl) gives the simplest ratio of ions.
• Metallic bonding: The electrostatic attraction between a lattice of positive metal ions and the surrounding sea of delocalised electrons. This attraction acts throughout the structure and is responsible for the characteristic properties of metals.

Crystal Structures: Ionic, Metallic & Molecular (Part 2/2)

• Delocalised electrons (in graphite): Electrons that are not associated with any single bond or atom but are spread across the entire layer of carbon atoms. Their mobility allows graphite to conduct electricity within each layer.

Giant Covalent Structures: Diamond & Graphite

• Giant covalent structure (macromolecular structure): A lattice of atoms bonded throughout by strong covalent bonds, forming one continuous network with no discrete molecules. The entire crystal is effectively a single giant molecule.

Energy Changes & States of Matter

• Intermolecular forces: Attractive forces between molecules — including London dispersion forces, permanent dipole–dipole interactions, and hydrogen bonding. These are the forces broken or formed during changes of state, not the covalent bonds within molecules.
• Enthalpy of fusion (ΔfusH): The enthalpy change when one mole of a solid melts to form a liquid at constant pressure. Always endothermic (Δ H > 0) because energy is required to partially overcome the intermolecular forces holding the solid lattice together.
• Enthalpy of vaporisation (ΔvapH): The enthalpy change when one mole of a liquid is converted to vapour at constant pressure. Always endothermic (Δ H > 0) because all remaining intermolecular forces between liquid molecules must be fully overcome.

Intermolecular Forces: Dipoles and Van der Waals

• Permanent dipole–dipole force: The electrostatic attraction between the δ+ end of one polar molecule and the δ− end of an adjacent polar molecule. These forces exist whenever molecules possess a permanent dipole — that is, whenever there is a net separation of charge across the molecule as a whole.
• Induced dipole–dipole force (London dispersion / van der Waals force): A temporary intermolecular attractive force present in all molecules. An instantaneous uneven distribution of electrons creates a temporary dipole in one molecule, which induces a dipole in an adjacent molecule. The resulting electrostatic attraction between the temporary and induced dipoles constitutes the force.

Hydrogen Bonding: Strength and Effects

• Hydrogen bond: A hydrogen bond forms between a hydrogen atom covalently bonded to a highly electronegative atom (N, O, or F) and a lone pair of electrons on a neighbouring electronegative atom (N, O, or F). It is represented by a dotted line in diagrams and is typically around 5–40 kJ mol^-1 in strength — much stronger than London dispersion forces, but considerably weaker than a covalent bond.

Enthalpy Change: Definitions

• Standard state: The most stable physical form of an element or compound at 298 K and 100 kPa. State symbols must reflect the standard state in any thermochemical equation.
• Standard enthalpy of combustion (Δc H^): The enthalpy change when one mole of a substance undergoes complete combustion in excess oxygen under standard conditions (298 K, 100 kPa), with all reactants and products in their standard states.
• Standard enthalpy of formation (Δf H^): The enthalpy change when one mole of a compound is formed from its constituent elements in their standard states under standard conditions (298 K, 100 kPa).

Enthalpy Changes

• Exothermic reaction: A reaction in which heat energy is transferred from the chemical system to the surroundings. The products have lower enthalpy than the reactants, so Δ H is negative.
• Endothermic reaction: A reaction in which heat energy is absorbed from the surroundings into the chemical system. The products have higher enthalpy than the reactants, so Δ H is positive.

Mean Bond Enthalpy: Definitions

• Mean Bond Enthalpy: The mean (average) enthalpy change when one mole of a specified type of covalent bond is broken homolytically in the gaseous phase, averaged over a range of different compounds containing that bond. Quoted in kJ mol⁻¹.

Calculating Enthalpy Changes Using Mean Bond Enthalpies

• Bond breaking is endothermic: Energy must be supplied to overcome the electrostatic attraction between bonded atoms. Breaking bonds always requires energy input, so the enthalpy change for bond breaking is always positive.
• Why mean bond enthalpies are averages: The C–H bond enthalpy in methane is not identical to the C–H bond enthalpy in ethane, propane, or any other compound. The precise bond energy depends on the electronic environment of the surrounding atoms. A mean bond enthalpy is calculated by averaging the bond dissociation enthalpies measured across many different compounds containing that bond type.

Hess's Law Calculations

• Hess's Law: The total enthalpy change for a chemical reaction is independent of the route taken, provided the initial and final conditions are the same.

Calorimetry Calculations: q=mcΔT and Molar Enthalpy Change

• Specific heat capacity (c): The energy required to raise the temperature of 1 gram of a substance by 1 °C. Units: J g⁻¹ °C⁻¹. For water and dilute aqueous solutions, c = 4.18 J g⁻¹ °C⁻¹ (provided in AQA exams).

Enthalpy Change Examples: Dissolution, Neutralization, Displacement, and Combustion

• q = mcΔT: The calorimetry equation relating heat transferred (q, in J) to the mass of solution (m, in g), the specific heat capacity (c, in J g⁻¹ °C⁻¹), and the temperature change (ΔT, in °C). To convert to kJ mol⁻¹, divide q by 1000 and then divide by the number of moles of the limiting reagent.

Temperature's Effect on Reaction Rate

• Rate of reaction: The change in concentration of a reactant or product per unit time. Expressed as mol dm⁻³ s⁻¹ when concentration is measured, or in other appropriate units (e.g. cm³ s⁻¹ for gas volume).

What are Catalysts?

• Catalyst: A substance that increases the rate of a chemical reaction without being changed in chemical composition or amount. The catalyst may participate in intermediate steps but is regenerated by the end of the reaction.

Maxwell-Boltzmann Distribution

• Most probable energy: The energy value at the peak of the Maxwell-Boltzmann curve. More molecules possess this energy than any other single energy value. It is always less than the mean energy because the distribution is positively skewed.

Collision Theory and Activation Energy

• Activation Energy (Ea): The minimum energy that colliding particles must possess for a collision to result in a chemical reaction. It represents the energy required to break the bonds in the reactants so that new bonds in the products can form. Measured in kJ mol^-1.

Concentration, Pressure, and Reaction Rates

• Collision frequency: The number of collisions between reacting particles per unit time. Increasing collision frequency increases the reaction rate, provided the proportion of successful collisions remains constant.

Kc Expressions and Calculations

• Square bracket notation [X]: The concentration of species X at equilibrium, expressed in mol dm⁻³. This notation is used throughout Kc expressions and must not be confused with the total amount of X in moles.

Factors Affecting Kc

• Equilibrium constant, Kc: The ratio of equilibrium concentrations of products to reactants, each raised to the power of their stoichiometric coefficient, for a reversible reaction at a specified temperature. Kc has a fixed value at a given temperature and changes only when temperature changes. For heterogeneous equilibria, the concentrations of pure solids and pure liquids are omitted from the Kc expression, as their concentrations are effectively constant.

Equilibrium: The Dynamic Balance

• Reversible reaction: A reaction in which the products can react together to re-form the reactants under appropriate conditions. Denoted by the symbol rather than a single-headed arrow.
• Dynamic equilibrium in a closed system: In a reversible reaction at equilibrium in a closed system, the forward and reverse reactions continue at equal rates and the concentrations of all species remain constant.
• Dynamic equilibrium: The state in a closed system where the rate of the forward reaction equals the rate of the reverse reaction. Concentrations of all species remain constant, but both reactions continue at the molecular level.

Le Chatelier's Principle: Shifting Equilibria

• Le Chatelier's Principle: If a system at dynamic equilibrium is subjected to a change in conditions, the position of equilibrium shifts in the direction that tends to oppose that change.
• Effect of a catalyst on equilibrium: A catalyst does not affect the position of equilibrium or the value of Kc. It provides an alternative reaction pathway with a lower activation energy, increasing the rates of both forward and reverse reactions equally, so equilibrium is reached more quickly but at the same position.

Industrial Equilibria: Finding the Balance

• Activation energy: The minimum energy that colliding particles must possess for a reaction to occur. Higher temperatures increase the proportion of particles with energy at or above this threshold, increasing reaction rate.

Oxidation States

• Oxidation state (formal definition): The charge an atom would have if the compound were fully ionic, assigned according to the priority rules above. Written as a signed integer with the sign before the number (e.g. +3, -2). Oxidation states of zero are written as 0, not +0.
• Oxidation and Reduction (in terms of oxidation state): Oxidation is an increase in oxidation state. Reduction is a decrease in oxidation state. The mnemonic OIL RIG (Oxidation Is Loss of electrons, Reduction Is Gain of electrons) connects this to electron transfer: losing electrons raises the oxidation state; gaining electrons lowers it.

Half-Equations and Redox

• Redox reaction: A reaction in which electron transfer occurs simultaneously: one species is oxidised (loses electrons) and another is reduced (gains electrons). The two processes cannot occur independently.
• Half-equation: An equation that represents either the oxidation or the reduction process in a redox reaction, with electrons shown explicitly. Oxidation half-equations show electrons as products; reduction half-equations show electrons as reactants.

Enthalpy Definitions

• Standard enthalpy of formation, Δf H°: The enthalpy change when one mole of a compound is formed from its constituent elements in their standard states under standard conditions (298 K, 100 kPa).
• Standard enthalpy of atomisation, Δat H°: The enthalpy change when one mole of gaseous atoms is formed from the element in its standard state under standard conditions.
• Ionisation energy: The enthalpy change when one mole of electrons is removed from one mole of gaseous atoms (first ionisation energy) or gaseous ions (successive ionisation energies) to form gaseous ions with one more positive charge, under standard conditions.
• Bond enthalpy (bond dissociation enthalpy): The enthalpy change when one mole of gaseous ions becomes hydrated by water molecules.
• Electron affinity: The enthalpy change when one mole of gaseous ions loses its hydration shell. It is almost always positive because energy is needed to overcome ion-dipole attractions.
• Lattice enthalpy of formation: The enthalpy change when one mole of an ionic compound is formed from its constituent gaseous ions under standard conditions. This process is exothermic (negative value) because strong electrostatic attractions are established.
• Lattice enthalpy of dissociation: The enthalpy change when one mole of an ionic compound is completely separated into its constituent gaseous ions under standard conditions. This process is endothermic (positive value) because all electrostatic attractions must be overcome.
• Enthalpy of hydration, Δhyd H°: The enthalpy change when one mole of gaseous ions is completely surrounded by water molecules to form an infinitely dilute aqueous solution under standard conditions.

Born-Haber Cycles and Lattice Enthalpy

• Lattice Enthalpy (Formation Convention): The enthalpy change when one mole of an ionic lattice is formed from its constituent gaseous ions under standard conditions (298 K, 100 kPa). It is always exothermic, so the value is always negative. Example: Na^+(g) + Cl^-(g) NaCl(s)
• Polarisation (in ionic compounds): The distortion of the electron cloud of an anion by the electric field of a neighbouring cation. Greater polarisation leads to greater sharing of electron density between the ions, introducing covalent character into what would otherwise be a purely ionic bond.

Enthalpy of Solution

• Enthalpy of solution, ΔsolH°: The enthalpy change when one mole of an ionic compound dissolves completely in sufficient water to produce an infinitely dilute solution, under standard conditions (298 K, 100 kPa). It can be exothermic (negative) or endothermic (positive) depending on the relative magnitudes of lattice enthalpy and hydration enthalpies.
• Enthalpy of hydration, ΔhydH°: The enthalpy change when one mole of a specified gaseous ion is completely hydrated by water molecules under standard conditions. Ion–dipole attractions form between the ion and water, releasing energy. Hydration enthalpies are always exothermic (negative). More highly charged and smaller ions have more exothermic hydration enthalpies.

Introduction to Gibbs Free Energy

• Enthalpy Change (Δ H): The heat energy exchanged between a system and its surroundings at constant pressure. Exothermic reactions have negative Δ H; endothermic reactions have positive Δ H. Alone, Δ H is not sufficient to determine whether a reaction is feasible.
• Entropy change (Δ S): Entropy change (Δ S) measures how the disorder or energy dispersal of a system changes during a process. A positive Δ S means the system becomes more disordered; a negative Δ S means it becomes more ordered. Units: J mol^-1 K^-1.

Gibbs Free Energy Calculations

• Gibbs Free Energy Change (Δ G): A thermodynamic quantity defined by Δ G = Δ H - TΔ S, where T is the absolute temperature in Kelvin. When Δ G < 0, the reaction is thermodynamically feasible (spontaneous). When Δ G > 0, the reaction is not feasible under those conditions. When Δ G = 0, the system is at the crossover point between feasible and non-feasible.

Introduction to Rate Equations

• Order of reaction (with respect to a reactant): The power to which the concentration of that reactant is raised in the experimentally determined rate equation. At A-level, orders are restricted to the values 0, 1, and 2.
• Rate constant (k): The proportionality constant in the rate equation rate = k[A]^m[B]^n. At a fixed temperature, k is constant for a given reaction. Its value increases with increasing temperature. Its units depend on the overall order of the reaction.

Rate Equation Calculations

• Rate equation: For a reaction involving reactants A and B, the rate equation takes the form: rate = k[A]^m[B]^n, where m is the order with respect to A, n is the order with respect to B, and k is the rate constant. The overall order is m + n.

The Arrhenius Equation

• Arrhenius constant, A: A pre-exponential factor with the same units as k. It represents the frequency of collisions with the correct orientation for reaction. Over moderate temperature ranges, A is treated as constant. It does not change with temperature in AQA calculations.

Concentration-Time Graphs and Initial Rates

• Instantaneous rate: The rate of reaction at one specific moment in time. It equals the gradient of the tangent drawn to the concentration-time curve at that moment. Units: mol dm⁻³ s⁻¹.
• Continuous monitoring method: An experimental technique in which a physical property related to concentration (such as absorbance, gas volume, or total mass) is measured at regular intervals throughout the reaction. The data are used to construct a concentration-time graph for the entire course of the reaction.

Reaction Mechanisms and Rate-Determining Step

• Rate-determining step (RDS): The slowest elementary step in a reaction mechanism. Only the species involved in the RDS (and any steps leading up to it) appear in the experimentally determined rate equation.

Kp and Partial Pressures

• Mole fraction (): The mole fraction of component A in a mixture is A = nAntotal, where nA is the moles of A and ntotal is the total moles of all gases present. Mole fractions are dimensionless and always sum to 1 across all components.
• Kp expression: For the general homogeneous gas-phase equilibrium aA(g) + bB(g) cC(g) + dD(g), the equilibrium constant Kp is defined as Kp = (pC)^c (pD)^d(pA)^a (pB)^b, where each p is the equilibrium partial pressure of that species.

Kp Calculations and Le Chatelier's Principle

• Units of Kp: The units of Kp depend on the overall change in moles of gas, Δn = (moles of gaseous products) − (moles of gaseous reactants). If Δn = 0, Kp is dimensionless. If Δn ≠ 0, the units are Pa^Δn or kPa^Δn, depending on the pressure unit used. Always derive units from the expression rather than memorising them.

Electrochemical Cells: Principles and Applications

• Half-cell: One half of an electrochemical cell, consisting of a metal electrode in contact with a solution of its ions (or a non-metal species in contact with an inert platinum electrode). Each half-cell hosts either oxidation or reduction.
• Non-rechargeable (irreversible) cell: A cell in which the electrode reactions are not reversible under normal operating conditions. Once the reactants are consumed, the cell cannot be restored and must be discarded. The alkaline dry cell is a common example.
• Rechargeable cell: A cell in which the electrode reactions are reversible. Passing an external electrical current through the cell in the reverse direction regenerates the original reactants, restoring the cell to its charged state. The lithium-ion cell is a widely used example.

Lithium Cells: Reactions and Function

• Cell EMF: The electromotive force of a cell, calculated as E°cell = E°cathode - E°anode. A positive value indicates the reaction is thermodynamically feasible under standard conditions.

Fuel Cells: Hydrogen-Oxygen and Societal Impact

• Fuel cell: An electrochemical device that generates an electric current by the continuous electrochemical oxidation of a fuel (hydrogen) at the anode and reduction of an oxidant (oxygen) at the cathode. Unlike a rechargeable battery, it does not store charge internally and does not require electrical recharging.

Electrode Potentials and Redox Reactions

• Standard hydrogen electrode (SHE): A half-cell consisting of hydrogen gas at 100 kPa bubbled over a platinum electrode immersed in a 1.00 mol dm^-3 solution of H^+(aq) at 298 K. The equilibrium established is 2H^+(aq) + 2e^- H2(g), and its electrode potential is defined as 0.00 V.

Electrochemical Cells: EMF and Representation

• IUPAC half-equation convention: All half-equations are written as reductions: oxidised form + ne^- reduced form. The electrode potential quoted in data tables always corresponds to this reduction direction.
• Conventional cell representation: Written as: anode half-cell | salt bridge | cathode half-cell. A single vertical line (|) represents a phase boundary between different states. A double vertical line (||) represents the salt bridge. The anode (oxidation) is always written on the left; the cathode (reduction) is always written on the right.

Practical: Measuring EMF of an Electrochemical Cell

• Electromotive force (EMF): The potential difference across a cell measured under zero-current conditions. It equals E°cathode - E°anode when both half-cells are at standard conditions. Measured in volts (V).

Titration Calculations and Curve Shapes

• Strong acid: An acid that is completely dissociated into ions in aqueous solution. For example, HCl(aq) → H⁺(aq) + Cl⁻(aq). The pH of a strong acid can be calculated directly from its concentration using pH = −log₁₀[H⁺].

Indicators and Weak/Strong Acid-Base Reactions

• Buffer region: The section of a weak acid–strong base (or weak base–strong acid) pH curve where pH changes only gradually as base is added, because a mixture of the weak acid and its conjugate base is present and resists pH change. It appears before the equivalence point.

Understanding Ka and Acid Dissociation

• Weak acid: A weak acid is one that only partially dissociates in aqueous solution, establishing an equilibrium between the undissociated acid HA and its ions H⁺ and A⁻. At any given moment, most molecules remain as HA.

Calculations with Ka and pKa

• Degree of dissociation: The fraction of acid molecules that have dissociated at equilibrium. For a weak acid, this is small — typically less than 5% — which underpins the approximation used in Ka calculations.
• pKa: pKa = -log10 Ka. A lower pKa means a stronger acid because the acid has a larger Ka.

Brønsted-Lowry Acids and Bases

• Brønsted-Lowry acid: A proton donor — a species that releases an H^+ ion to another species. In the reaction above, HCl donates its proton and is therefore the acid.
• Brønsted-Lowry base: A proton acceptor — a species that uses a lone pair of electrons to bond with an H^+ ion donated by an acid. In the reaction above, NH3 accepts the proton and is therefore the base.

Understanding and Calculating pH

• pH scale: A logarithmic scale used to express hydrogen ion concentration in aqueous solution. Defined by pH = -log10[H+], where [H+] is in mol dm-3. Each unit change in pH corresponds to a tenfold change in [H+].

Kw Calculations and Temperature Dependence

• Ionic product of water, Kw: Kw = [H+][OH-], where both concentrations are in mol dm⁻³. At 298 K, Kw = 1.00 × 10⁻¹⁴ mol² dm⁻⁶. This relationship holds in any aqueous solution — acidic, basic, or neutral — at a given temperature.

Inorganic chemistry

36 cards from 23 lessons. Each term links to the lesson it comes from.

s, p, d, f Block Elements

• s block element: An element whose highest-energy electrons occupy an s subshell. Found in Groups 1 and 2 of the Periodic Table. The s subshell holds a maximum of 2 electrons.
• d block element: An element whose highest-energy electrons occupy a d subshell. Found in Groups 3 to 12 of the Periodic Table. The d subshell holds a maximum of 10 electrons, giving the d block its characteristic width of ten groups.

Period 3 Trends: Atomic Radius and Ionisation Energy

• Effective nuclear charge: The net positive charge felt by an outer electron. It equals the actual nuclear charge minus the shielding effect of inner electrons. Across Period 3, shielding stays roughly constant while nuclear charge rises, so effective nuclear charge increases from Na to Ar.

Group 2 Trends: Atomic Properties and Reactions

• Shielding (screening): The reduction in the attractive force experienced by outer electrons due to repulsion from inner electron shells. Greater shielding means outer electrons are held less tightly by the nucleus.

Group 2 Hydroxides: Solubility and Uses

• Sparingly soluble: A substance that dissolves only to a very small extent in water. For Mg(OH)2, this means only a tiny amount dissociates to give OH- ions in solution, producing a mildly alkaline pH rather than a strongly alkaline one.

Group 2 Sulfates: Solubility, Testing, and Applications

• Positive result for sulphate ions: A white precipitate forms when acidified barium chloride solution is added to a solution containing sulphate ions. The precipitate is barium sulphate, BaSO4(s), which is insoluble in dilute hydrochloric acid.

Chlorine Reactions with Water

• Disproportionation: A redox reaction in which the same element is simultaneously oxidised and reduced. In the reaction of chlorine with water, chlorine starts at oxidation state 0 in Cl2. One chlorine atom is reduced to Cl^- (oxidation state -1) and the other is oxidised to ClO^- (oxidation state +1). Both changes happen in the same reaction.

Chlorine in Water Treatment

• Reaction of chlorine with water: Chlorine dissolves in water to form a mixture of hydrochloric acid and chloric(I) acid (hypochlorous acid): Cl2(aq) + H2O(l) HCl(aq) + HOCl(aq). The equilibrium lies to the left, so the reaction is only partial.
• Reaction of chlorine with cold dilute aqueous sodium hydroxide: Chlorine reacts with cold dilute NaOH(aq) to form sodium chloride, sodium chlorate(I) (sodium hypochlorite), and water: Cl2(aq) + 2NaOH(aq) NaCl(aq) + NaOCl(aq) + H2O(l). This is also a disproportionation reaction.

Halogen Trends: Electronegativity, Boiling Point & Oxidising Ability

• Shielding effect: The reduction in the attractive force experienced by outer electrons due to repulsion from inner-shell electrons. More inner shells means greater shielding, reducing the effective nuclear charge felt by bonding electrons.
• Oxidising agent: A species that accepts electrons from another species, causing that other species to be oxidised. In Group 7, a halogen acts as an oxidising agent when it gains an electron to form a halide ion. The stronger the oxidising agent, the more readily it accepts electrons.

Halide Ions: Reducing Ability & Sulfuric Acid Reactions

• Polarisability and electron donation: As you descend Group 7, each halide ion has more electron shells and a larger ionic radius. The outer electrons are further from the nucleus and are shielded by more inner electron shells. This means the outer electrons are held less tightly and can be donated more readily — the ion is more polarisable and therefore a stronger reducing agent.

Identifying Halide Ions: Silver Nitrate & Ammonia Tests

• Acidification with dilute nitric acid: Adding dilute nitric acid (HNO3) to the test solution before silver nitrate is introduced. This removes interfering anions — particularly carbonate (CO3^2-), sulphite (SO3^2-), and sulphate (SO4^2-) — which would otherwise react with silver ions to form white or pale precipitates, producing a false positive result for chloride. Nitric acid is the only suitable acid because nitrate ions do not react with silver ions; using hydrochloric acid would introduce chloride ions, and using sulphuric acid would introduce sulphate ions, both of which would interfere with the test.
• Diamminesilver(I) complex ion: The complex ion [Ag(NH3)2]^+ formed when silver ions react with ammonia molecules. When a silver halide dissolves in ammonia, the silver ions are removed from the equilibrium by forming this stable complex, driving the dissolution of the precipitate.

Period 3 Elements and Reactions with Oxygen and Water

• Vigorous reaction: In the context of Period 3 metals with water, 'vigorous' means rapid effervescence, heat generation, and visible movement of the metal on the water surface. Sodium is the clearest example in Period 3.
• Amphoteric oxide: An oxide that reacts with both acids and bases to produce a salt and water. Aluminium oxide (Al2O3) is the Period 3 example. It does not react with water under normal conditions.

Structure, Bonding, and Melting Points of Period 3 Oxides

• Van der Waals forces (London dispersion forces): Weak, temporary intermolecular forces arising from instantaneous dipoles in electron clouds. They are present between all simple molecules, including P4O10 and SO3. Note that SO3 is a polar molecule (trigonal planar with polar S=O bonds giving a net dipole), so dipole-dipole forces also act between SO3 molecules in addition to van der Waals forces. Because all these intermolecular forces are weak, little energy is needed to overcome them, giving low melting points.

Octahedral and Tetrahedral Complex Ions

• Octahedral complex: A complex ion in which six ligands are arranged around a central metal ion at the vertices of a regular octahedron. Adjacent bond angles are 90° and trans ligands are 180°. Coordination number = 6. Typical ligands: H2O and NH3.
• Tetrahedral complex: A complex ion in which four ligands are arranged around a central metal ion at the vertices of a tetrahedron. Bond angles are 109.5°. Coordination number = 4. Typical ligand: Cl^-.

Isomerism in Complex Ions

• Cis isomer (octahedral): The isomer in which two identical monodentate ligands occupy adjacent coordination sites (bond angle of 90°) in an octahedral complex.
• [Ag(NH3)2]+: The diamminesilver(I) ion — a linear complex formed when excess ammonia is added to silver nitrate solution. It is the active species in Tollens' reagent, used to distinguish aldehydes from ketones.

Monodentate Ligand Exchange

• Monodentate ligand: A ligand that forms one co-ordinate bond to the central metal ion by donating a single lone pair. H2O donates via oxygen, NH3 donates via nitrogen, and Cl- donates via one of its four lone pairs.

Multidentate Ligands and the Chelate Effect

• Ethane-1,2-diamine (en): H2NCH2CH2NH2 — a bidentate ligand in which both nitrogen atoms each carry a lone pair. The two-carbon chain between them holds the nitrogen atoms at exactly the right distance to bridge across one metal ion, forming a stable five-membered chelate ring.
• EDTA4–: A hexadentate ligand with the formula [CH2N(CH2COO–)2]2 that wraps around a metal ion using two nitrogen atoms and four oxygen atoms, forming six coordinate bonds and occupying all six coordination sites of an octahedral metal complex in a single ligand.

Haemoglobin and Ligand Substitution

• Haem: An iron(II) complex in which a porphyrin ring (a large, flat, cyclic organic molecule) acts as a multidentate ligand, co-ordinating to Fe(II) through four nitrogen atoms. The complex sits within the haemoglobin protein and is the site of oxygen binding.
• Ligand substitution (in the context of CO toxicity): The replacement of an oxygen ligand co-ordinately bonded to Fe(II) in haemoglobin by a carbon monoxide ligand. Because CO binds far more strongly than O2, this substitution is not readily reversed, preventing oxygen transport and causing toxicity.

Introduction to Variable Oxidation States

• Variable oxidation states: The capacity of a transition metal to exist in more than one stable oxidation state, resulting from the similar energies of the 3d and 4s subshells. Electrons can be removed from either subshell, or in different numbers, without an insurmountable energy cost.
• Tollens' reagent: An alkaline solution of the diamminesilver(I) complex ion [Ag(NH3)2]^+, prepared by adding dilute ammonia to silver nitrate solution until the initial precipitate of silver oxide just dissolves. It is used to distinguish aldehydes from ketones.

Intro to Transition Metals & Complex Formation

• Transition metal (AQA definition): An element that forms at least one stable ion with an incomplete d sub-level. This distinguishes true transition metals from other d-block elements such as scandium and zinc.
• Ligand: A molecule or ion that forms a co-ordinate bond with a transition metal by donating a lone pair of electrons to the central metal atom or ion. Common examples include water (H₂O), ammonia (NH₃), chloride ions (Cl⁻), and cyanide ions (CN⁻).
• Co-ordination number: The number of co-ordinate bonds formed between the central metal atom or ion and its surrounding ligands in a complex. It is determined by counting bonds, not ligands — a bidentate ligand contributes 2 to the co-ordination number.

Coloured Ions, Variable Oxidation States & Catalysis

• d-d transition: A transition metal is an element that forms at least one stable ion with a partially filled d-subshell.

Metal Aqua Ions and Acidity

• [M(H₂O)₆]²⁺ ions (AQA specification): Formed by Fe²⁺ and Cu²⁺ in aqueous solution. Each metal ion is surrounded by six water ligands in an octahedral arrangement, coordinated through the oxygen lone pairs.

Reactions of Metal Ions with Bases

• Precipitation reaction (metal ion + OH–): The metal cation in aqueous solution reacts with hydroxide ions to form an insoluble metal hydroxide. General form: M^n+(aq) + n OH–(aq) → M(OH)n(s). The precipitate forms immediately on addition of the base.
• Amphoteric character of Al(OH)3: Al(OH)3 reacts with both acids and bases. With excess NaOH: Al(OH)3(s) + OH–(aq) → [Al(OH)4]–(aq). With dilute acid (e.g. HCl): Al(OH)3(s) + 3H+(aq) → Al3+(aq) + 3H2O(l). The precipitate dissolves in both cases.

Practical: Identifying Metal Ions in Solution

• Aqua complex: A complex ion in which water molecules act as ligands surrounding a central metal ion, written as [M(H2O)6]^n+. The colour of the aqua complex is characteristic of the metal ion and its oxidation state.

Organic chemistry

102 cards from 68 lessons. Each term links to the lesson it comes from.

Understanding Reaction Mechanisms

• Homolytic fission: Each atom receives one electron from the shared pair when the bond breaks. This produces two species, each with one unpaired electron — these are called free radicals. The unpaired electron is represented by a dot placed next to the atomic symbol, for example Cl· or CH₃·.
• Heterolytic fission: Both electrons from the shared pair go to one atom when the bond breaks. This produces one positively charged ion (cation) and one negatively charged ion (anion). The atom that takes both electrons becomes the anion; the atom that loses both becomes the cation.

Curly Arrows: Bond Formation and Breaking

• Bond formation via lone pair: When a nucleophile donates a lone pair to form a new covalent bond, the curly arrow starts from the lone pair on the donor atom and points towards the atom that will receive the bond. This is the basis of all nucleophilic attack steps.
• Bond breaking (heterolytic): When a covalent bond breaks heterolytically, both electrons move together to one atom. The curly arrow starts from the bond and points to the atom that receives both electrons, generating a lone pair on that atom and a positive charge on the atom left behind.

Formulae in Organic Chemistry

• Empirical formula: The simplest whole-number ratio of atoms of each element present in a compound. It does not necessarily reflect the actual number of atoms in one molecule.
• Molecular formula: The actual number of atoms of each element present in one molecule of a compound. It is always a whole-number multiple of the empirical formula.
• General formula: An algebraic expression that represents the molecular formula of any member of a homologous series. The variable n represents the number of carbon atoms (or repeat units) in the chain.
• Structural formula: A representation that shows the arrangement of atoms and bonds in a molecule, written in a condensed linear form. Each carbon and its attached hydrogens are grouped together, and functional groups are written out explicitly.
• Displayed formula: A two-dimensional representation in which every atom is shown as its symbol and every covalent bond is shown as a line. No atoms or bonds are implied or omitted.
• Skeletal formula: A simplified structural representation in which the carbon skeleton is drawn as a zig-zag of lines. Each line junction or end-point represents a carbon atom. Hydrogen atoms bonded to carbon are omitted entirely; only functional groups and heteroatoms are shown explicitly.

IUPAC Nomenclature and Homologous Series

• Carbon-chain stem: The part of an IUPAC name that encodes the number of carbon atoms in the longest continuous carbon chain. The stems are: meth- (1C), eth- (2C), prop- (3C), but- (4C), pent- (5C), hex- (6C).

Structural Isomerism: Types and Drawing

• Structural isomers: Compounds that have the same molecular formula but different structural formulae. The atoms are connected in a different arrangement, giving each isomer distinct physical and chemical properties.

Stereoisomerism and E-Z Isomers

• Stereoisomers: Molecules with the same molecular formula and the same sequence of bonded atoms, but with a different arrangement of atoms in space.

Cracking Alkanes: An Introduction

• Free radical: A species with an unpaired electron, represented by a dot (•). In cracking, homolytic fission of C–C bonds generates carbon-centred free radicals that drive the chain of fragmentation reactions.
• Zeolite: An aluminosilicate material with a highly porous, regular structure that provides a large surface area. Used as the catalyst in catalytic cracking; its acidic sites promote C–C bond breaking at lower temperatures than thermal cracking requires.

Thermal vs. Catalytic Cracking

• Thermal cracking: A process in which long-chain alkanes are broken down by high temperature and high pressure, producing a high percentage of alkenes alongside shorter-chain alkanes. No catalyst is used.
• Catalytic cracking: A process in which long-chain alkanes are broken down at high temperature, slight pressure, and in the presence of a zeolite catalyst. It is used mainly to produce motor fuels and aromatic hydrocarbons.

Introduction to Alkanes and Fractional Distillation

• Saturated hydrocarbon: A hydrocarbon in which all carbon-carbon bonds are single bonds. Every carbon atom uses its four available bonding positions to form single bonds only, either to other carbon atoms or to hydrogen atoms.

Fuels and Combustion

• Complete combustion: The reaction of a hydrocarbon with excess oxygen, producing carbon dioxide and water as the only carbon- and hydrogen-containing products. The general equation for an alkane CnH2n+2 is: CnH2n+2 + (3n+1/2)O2 nCO2 + (n+1)H2O
• Sulphur dioxide formation from combustion: When a hydrocarbon fuel containing sulphur impurities is burned, the sulphur is oxidised to sulphur dioxide: S(s) + O2(g) SO2(g). This reaction occurs alongside the combustion of the hydrocarbon itself.

Pollutants and Catalytic Converters

• NOx (Nitrogen Oxides): A mixture of nitrogen monoxide (NO) and nitrogen dioxide (NO2) formed when atmospheric nitrogen and oxygen react at the high temperatures inside a combustion engine. NOx contributes to acid rain (via conversion to nitric acid in the atmosphere) and to photochemical smog.
• Heterogeneous catalysis (in a catalytic converter): The catalyst (solid platinum-rhodium alloy) is in a different phase from the reactants (gaseous exhaust pollutants). Reactant molecules adsorb onto the catalyst surface, react, and then desorb as less harmful products. The catalyst is not consumed in the process.

Sulfur Dioxide Removal

• Flue gas: The mixture of gases produced by combustion in a furnace or power station boiler, which travels up the chimney (flue) before being released. Contains CO2, water vapour, N2, and pollutants such as SO2.

Elimination Reactions: Nucleophiles vs. Bases

• The dual role of OH⁻: When OH⁻ attacks the carbon bearing the halogen directly, it acts as a nucleophile and substitution occurs. When OH⁻ abstracts a proton from an adjacent carbon, it acts as a base and elimination occurs. Both pathways compete simultaneously.

Understanding Nucleophilic Substitution

• Polar C–X bond in halogenoalkanes: The carbon–halogen bond (C–X) is polar because the halogen is more electronegative than carbon. This creates a δ+ charge on carbon and a δ− charge on the halogen, making the carbon susceptible to attack by nucleophiles.

Mechanisms and Reaction Rates

• Leaving group: The atom or group that departs with the bonding electron pair during a substitution reaction. In halogenoalkanes, the halide ion (X^-) is the leaving group. A good leaving group is a stable anion — iodide is the best, fluoride the worst.
• Activation energy and bond enthalpy link: In nucleophilic substitution, breaking the C-X bond contributes significantly to the activation energy of the reaction. A lower C-X bond enthalpy means less energy is needed to reach the transition state (S_N2) or to form the carbocation (S_N1), so the reaction rate increases.

Ozone Layer and CFCs

• Ozone (O3): A triatomic allotrope of oxygen found naturally in the stratosphere. It is beneficial because it absorbs ultraviolet radiation, preventing the majority of harmful UV from reaching the Earth's surface.

CFCs and Legislation

• Chlorine radical: A chlorine atom with an unpaired electron, written as Cl. Generated when UV radiation breaks the C–Cl bond in a CFC molecule in the stratosphere. Acts as a catalyst in the chain destruction of ozone.

Electrophilic Addition Mechanisms

• Electrophile: A species that accepts a pair of electrons from an electron-rich site. Electrophiles are either positively charged (e.g. H+) or have a partial positive charge on an atom (e.g. the H in H-Br, or the Br in Br-Br when polarised by the pi bond).
• Bromine water test for unsaturation: A qualitative test in which bromine water (orange-brown) is shaken with an unknown compound. If a C=C double bond is present, the bromine adds across the pi bond via electrophilic addition, and the solution is decolourised to colourless. A saturated compound (no C=C) does not react under these conditions, and the orange-brown colour persists.

Carbocation Stability and Product Formation

• Carbocation intermediate: A carbon atom carrying a full positive charge, bonded to only three other atoms, with an empty p orbital perpendicular to the plane of its three bonds. It forms when the pi bond of an alkene donates both electrons to an incoming electrophile, leaving one carbon electron-deficient.

Addition Polymer Structure and Naming

• Addition polymerisation: A reaction in which many monomer molecules, each containing a C=C double bond, join together to form a long polymer chain. No other product is formed. The double bond in each monomer opens to create two new single bonds linking adjacent monomer units.

Properties, Uses, and Historical Context of Addition Polymers

• Chemical unreactivity of addition polymers: Addition polymers are unreactive because their backbone contains only C-C and C-H single (sigma) bonds. These bonds are strong and non-polar (or only weakly polar), so they do not attract electrophiles, nucleophiles, or other reagents under normal conditions. There are no C=C double bonds remaining to undergo addition reactions.

Alkenes: Structure and Bonding

• Alkene: An unsaturated hydrocarbon containing at least one carbon-carbon double covalent bond (C=C). The general formula for a straight-chain alkene is CnH2n.
• Sigma bond (σ bond): A covalent bond formed by the direct, end-on overlap of two atomic orbitals along the internuclear axis. It is cylindrically symmetrical about the bond axis and allows rotation. Every single bond is a sigma bond; the first bond in a double bond is also a sigma bond.

Biofuels and Fermentation

• Biofuel: A fuel produced from recently living biological material (biomass). Because the carbon in biomass was recently fixed from atmospheric CO2 by photosynthesis, biofuels are considered potentially carbon neutral — unlike fossil fuels, which release carbon that has been locked away for millions of years.

Alkene Hydration and Environmental Impact

• Electrophilic addition: A reaction in which an electrophile attacks the pi bond of an alkene. The pi bond breaks heterolytically, forming a carbocation intermediate. A nucleophile then attacks the carbocation to complete the addition.

Alcohol Elimination Reactions: Mechanisms and Alkene Formation

• Acid-catalysed elimination: The removal of water (H and OH) from an alcohol in the presence of a concentrated acid catalyst and heat, producing an alkene. The acid is regenerated at the end of the mechanism.
• Zaitsev's rule: The major product of an elimination reaction is the most substituted alkene — that is, the alkene with the greatest number of alkyl groups attached to the double-bond carbons. More substituted alkenes are more stable due to hyperconjugation and inductive electron donation from alkyl groups.

Practical Skills: Elimination and Distillation

• Elimination (dehydration of an alcohol): The removal of water (H2O) from an alcohol molecule to form an alkene. The OH group and an adjacent H atom are removed. Concentrated acid (e.g. H3PO4) at around 170 °C catalyses this process. Alternatively, passing alcohol vapour over hot aluminium oxide (Al2O3) at approximately 350 °C achieves the same result.

Alcohol Oxidation: Types and Equations

• Primary alcohol: A primary alcohol has the carbon bonded to the -OH group attached to either no other carbon atoms (as in methanol) or to only one other carbon atom. Examples include methanol, ethanol, and propan-1-ol.

Controlling Oxidation and Identifying Carbonyls

• Distillation (in this context): The technique of heating the reaction mixture and collecting the volatile product (the aldehyde) as it forms, preventing further oxidation to the carboxylic acid.

Understanding IR Spectroscopy

• Wavenumber: The number of wave cycles per centimetre, measured in cm⁻¹. On an IR spectrum, the x-axis runs from around 4000 cm⁻¹ on the left to 500 cm⁻¹ on the right. A bond absorbing at a given wavenumber appears as a dip (trough) in the spectrum, since that radiation is absorbed rather than transmitted.

Identifying Molecules and Global Warming

• Bond vibration (IR absorption condition): A bond absorbs infrared radiation only if the vibration causes a change in the dipole moment of the molecule. Symmetrical, non-polar bonds such as O₂ and N₂ do not absorb IR radiation because their vibrations produce no dipole change.
• Enhanced greenhouse effect: The intensification of the natural greenhouse effect caused by increased atmospheric concentrations of greenhouse gases — principally CO₂ from combustion of fossil fuels and CH₄ from agriculture and decomposition — leading to additional warming of the Earth’s surface beyond the natural baseline.

Using Mass Spectrometry to Determine Molecular Formula

• Molecular ion (M⁺): The ion formed when a molecule loses one electron during ionisation in the mass spectrometer, giving a species with a charge of +1. Its m/z value (where z = 1) equals the relative molecular mass of the compound to the precision of the instrument.

Functional Group Identification

• Tollens' reagent (ammoniacal silver nitrate): A solution of silver nitrate dissolved in aqueous ammonia, giving the complex ion [Ag(NH3)2]^+. When warmed with an aldehyde, the silver(I) ions are reduced to silver metal, depositing a silver mirror on the inside of the test tube.

Drawing and Understanding Enantiomers

• Chiral centre: A carbon atom carrying four different substituents, producing a molecule that is non-superimposable on its mirror image. Denoted with an asterisk (*) in structural formulae.
• Racemic mixture (racemate): An equimolar mixture of both enantiomers of a chiral compound. A racemic mixture has no net effect on plane-polarised light because the rotations of the two enantiomers cancel exactly.

Redox Reactions of Aldehydes and Ketones

• Carbonyl compound: An organic compound containing the C=O functional group. Aldehydes have the carbonyl at the terminal position (RCHO); ketones have it flanked by two carbon substituents (RCOR').
• Oxidation of an aldehyde: Aldehydes are readily oxidised to carboxylic acids. The CHO group gains an oxygen atom (or loses two hydrogen atoms) to form COOH. In overall equation terms, this can be written as RCHO + [O] → RCOOH, where [O] represents the oxygen from the oxidising agent.

Hydroxynitriles from Carbonyls

• Racemic mixture: An equal mixture of both enantiomers of a chiral compound, produced when a reaction creates a chiral centre from a planar (achiral) starting material with no chiral influence. A racemic mixture does not rotate plane-polarised light overall because the rotations of the two enantiomers cancel.

Acyl Compounds: Structure and Properties

• Acyl chloride structure: General formula RCOCl. The carbonyl carbon is bonded to an alkyl or aryl group R, a double-bonded oxygen, and a chlorine atom. Example: ethanoyl chloride, CH3COCl.
• Acid anhydride structure: General formula (RCO)2O. Two acyl groups are linked through a central oxygen atom, giving the -CO-O-CO- linkage. Example: ethanoic anhydride, (CH3CO)2O.
• Amide structure: General formula RCONH2 (primary amide). The carbonyl carbon is bonded to a nitrogen atom that carries one or two hydrogen atoms (or alkyl groups in secondary and tertiary amides). Example: ethanamide, CH3CONH2.
• Melting point as a purity test for organic solids: A pure organic solid melts at a precise, sharp temperature (within 1-2 °C) that matches the literature value. An impure sample melts over a wider range and at a lower temperature than expected. The width of the melting range is an indicator of purity: the narrower the range, the purer the sample.
• Boiling point as a purity test for organic liquids: A pure organic liquid boils at a precise, constant temperature under a given pressure. An impure liquid boils over a range of temperatures. Distillation is used to purify the liquid, and the fraction collected at the expected boiling point is taken as the pure compound.

Acylation Reactions: Mechanisms and Applications

• Electrophilic carbonyl carbon: In acyl chlorides (RCOCl), the carbonyl carbon is made even more electron-deficient because the electronegative chlorine atom withdraws further electron density by induction. This makes acyl chlorides more reactive acylating agents than acid anhydrides, which have a less electron-withdrawing leaving group.

Practical: Preparing and Purifying an Ester

• Acid catalyst in esterification: Concentrated sulphuric acid donates a proton to the carbonyl oxygen of the carboxylic acid, activating it towards nucleophilic attack by the alcohol. The catalyst lowers activation energy and speeds up both the forward and reverse reactions equally, so it does not shift the equilibrium position — it only increases the rate at which equilibrium is reached.
• Anhydrous drying agent: An anhydrous (water-free) inorganic salt that absorbs water from an organic liquid by forming a stable hydrate. Common examples include anhydrous calcium chloride (CaCl2), anhydrous magnesium sulphate (MgSO4), and anhydrous sodium sulphate (Na2SO4). The drying agent is added in small portions until it no longer clumps together, indicating that all free water has been absorbed. Note: anhydrous calcium chloride cannot be used to dry alcohols because it forms a complex with them (e.g. CaCl2 × 4C2H5OH); it is suitable here only because ethanol has been removed during the washing stage. Anhydrous magnesium sulphate or sodium sulphate are preferred when alcohol may still be present.

Carboxylic Acids: Structure, Properties, and Formation

• Carboxyl group: The functional group -COOH, consisting of a carbonyl (C=O) and a hydroxyl (O-H) bonded to the same carbon. It is the defining feature of all carboxylic acids.

Esters: Formation, Uses, and Hydrolysis

• Ester functional group: The -COO- group, in which a carbonyl carbon (C=O) is bonded to an oxygen atom that is itself bonded to a carbon chain. This group is the defining feature of all esters.

Fats, Oils, Soaps and Biodiesel

• Propane-1,2,3-triol (glycerol): A triol with the molecular formula C3H5(OH)3. It has hydroxyl groups on each of its three carbon atoms, allowing it to form three ester bonds with carboxylic acids to produce a triglyceride.
• Saponification: The alkaline hydrolysis of a triglyceride. A hot concentrated alkali (NaOH or KOH) cleaves all three ester bonds, producing glycerol and the metal salts of long-chain carboxylic acids (soap). The reaction is irreversible under these conditions.

Benzene's Structure and Stability

• Intermediate bond length: In benzene, all six C–C bonds have a length of approximately 0.140 nm, which is between the length of a C–C single bond (~0.154 nm) and a C=C double bond (~0.134 nm). This is direct evidence that the bonding cannot be described as alternating single and double bonds.

Benzene Reactions: Substitution vs. Addition

• Delocalisation energy (stabilisation energy): The extra stability benzene possesses relative to a hypothetical cyclohexatriene with three isolated double bonds. Experimentally determined by comparing actual and theoretical enthalpies of hydrogenation: benzene's actual value is approximately 208 kJ mol^-1 less exothermic than the theoretical value of 360 kJ mol^-1, giving a delocalisation energy of approximately 152 kJ mol^-1.

Friedel-Crafts Acylation

• Electrophilic Aromatic Substitution (EAS): A reaction in which an electrophile replaces a hydrogen atom on an aromatic ring. The delocalised π system attacks the electrophile, forming a Wheland intermediate, which then loses a proton to restore aromaticity. Substitution is always the outcome — not addition — because addition would permanently destroy the stable delocalised ring.
• Acylium Ion: The electrophile in Friedel-Crafts acylation, formed by the reaction of an acyl chloride with AlCl3. The aluminium chloride accepts the chloride ion, leaving a positively charged species of the form RCO^+, stabilised by resonance between R-+C=O and R-C+O.

Amine Base Strength

• Weak base: A base that only partially accepts protons in aqueous solution, establishing an equilibrium in which the majority of the base remains in its unprotonated form. Amines are weak bases because the equilibrium for proton acceptance lies to the left.

Amines as Nucleophiles: Substitution Reactions

• Nucleophile (exam-standard definition): A species that donates an electron pair to an electron-deficient (electrophilic) atom, forming a new covalent bond. Ammonia and amines are nucleophiles because the nitrogen atom possesses a lone pair of electrons.

Quaternary Ammonium Salts as Cationic Surfactants

• Quaternary ammonium ion: A cation of the form [NR4]^+, where each R is an alkyl or aryl group. The nitrogen has no lone pair and cannot act as a base or nucleophile. The positive charge is permanent and does not depend on pH.

Preparation of Amines

• Excess ammonia (in amine synthesis): Using a large molar excess of ammonia relative to the halogenoalkane drives the equilibrium towards the primary amine product and suppresses further alkylation. Without excess, secondary and tertiary amines — and quaternary ammonium salts — form as by-products.

Condensation Polymerisation: Monomers and Repeating Units

• Condensation polymerisation: A polymerisation reaction in which monomers with two functional groups react repeatedly, eliminating a small molecule (typically water) at each bond-forming step to produce a long-chain polymer with characteristic linkage groups.
• Dicarboxylic acid: A molecule carrying two carboxylic acid groups (–COOH), one at each end of the carbon chain. Examples include hexanedioic acid (adipic acid), used in nylon 6,6 synthesis.

Properties, Intermolecular Forces and Uses of Condensation Polymers

• Permanent dipole-dipole interaction: An attractive intermolecular force between the permanent partial charges (δ+ and δ-) on neighbouring polar molecules or polar sections of polymer chains. Stronger than London forces but weaker than hydrogen bonds.

Biodegradability of Polymers

• Chemical inertness (polyalkenes): Polyalkenes are described as chemically inert because their backbones contain only C-C and C-H bonds. These bonds are non-polar, strong, and do not react with water, acids, or alkalis under normal environmental conditions.
• Ester linkage: The functional group -COO- formed when a carboxylic acid reacts with an alcohol. In polyesters, this linkage repeats throughout the backbone and can be cleaved by hydrolysis under acidic or alkaline conditions.

Polymer Disposal and Recycling

• Landfill: A method of solid waste disposal in which materials are buried in the ground, typically in lined pits designed to prevent leachate from contaminating groundwater.

DNA Structure and Base Pairing

• Nucleotide: A nucleotide consists of three components bonded together in a specific sequence: a phosphate ion, bonded to 2-deoxyribose (a pentose sugar), which is in turn bonded to one of the four nitrogenous bases — adenine (A), thymine (T), guanine (G), or cytosine (C). The phosphate–sugar bond and the sugar–base bond are both covalent.

DNA Polymerization

• 2-Deoxyribose: The pentose sugar in every DNA nucleotide. It has five carbons, numbered 1' to 5'. The critical feature is the absence of a hydroxyl (–OH) group at the 2' carbon — there is only a hydrogen atom there instead. This distinguishes DNA from RNA (which uses ribose, with –OH at the 2' position) and contributes to DNA's greater chemical stability.
• Phosphodiester bond: A covalent bond in which a phosphate group is esterified to two sugar molecules simultaneously — one via the 5' carbon and one via the 3' carbon of adjacent deoxyribose units. The bond is formed by a condensation reaction with the elimination of water. It is the bond that links successive nucleotides in the DNA backbone.

Amino Acids and Zwitterions

• General structure of an alpha-amino acid: An alpha-amino acid has the formula H2N-CH(R)-COOH, where R is the side chain that distinguishes one amino acid from another. The amino group and carboxyl group are both bonded to the alpha carbon.
• Zwitterion: A dipolar ion formed when the -COOH group of an amino acid donates a proton to the -NH2 group within the same molecule, producing -COO^- and -NH3^+ simultaneously. Net charge = zero.

Enzymes: Stereospecificity and Catalysis

• Active site: A region on an enzyme, formed by the precise folding of the protein chain, with a complementary shape and chemical environment to a specific substrate. Temporary bonds — including hydrogen bonds, ionic interactions, and van der Waals forces — form between the active site residues and the substrate during catalysis.
• Stereospecificity: The ability of an enzyme active site to bind to only one enantiomeric form of a chiral substrate or drug. This arises because the active site presents functional groups in a fixed three-dimensional arrangement, so only the enantiomer with the correct spatial configuration can form the necessary interactions simultaneously.

Enzyme Inhibition and Drug Design

• Enzyme inhibitor (drug context): A drug molecule designed with a shape and functional group arrangement complementary to a target enzyme's active site, so that it binds there and prevents the substrate from binding, thereby reducing or eliminating the enzyme's catalytic activity.

Cisplatin: How it Works

• cis isomer: In cisplatin, the two chloride ligands are positioned on the same side of the square planar complex (cis arrangement). The trans isomer, with chloride ligands on opposite sides, is pharmacologically inactive — the geometry is critical to the drug's function.
• Pt–N coordinate bond with guanine: The bond formed when the N7 nitrogen atom of a guanine base donates its lone pair to the Pt(II) centre in activated cisplatin, displacing the water ligand. This is the key bond responsible for cisplatin's anticancer activity.

Cisplatin: Balancing Benefits and Risks

• Hydrolysis of cisplatin: The substitution of one or both chloride ligands by water molecules inside the cell, producing [Pt(NH3)2(H2O)Cl]+ and then [Pt(NH3)2(H2O)2]2+. These species are electrophilic and react readily with nucleophilic sites on DNA.

Peptide Structures and Hydrolysis

• General amino acid structure: H2N-CH(R)-COOH, where R is the side chain. At physiological pH, amino acids exist as zwitterions: +H3N-CH(R)-COO-. For AQA drawing questions, use the un-ionised form H2N-CH(R)-COOH unless the question specifies otherwise.
• Hydrolysis of a peptide link: The reaction of a peptide bond (-CO-NH-) with water, breaking the C-N bond and adding -OH to the carbonyl carbon and -H to the nitrogen. This regenerates the carboxyl group (-COOH) and amino group (-NH2) of the constituent amino acids. Conditions: dilute acid or alkali, heat; or enzyme at body temperature.

Green Chemistry Principles in Organic Synthesis

• Non-hazardous starting materials: Reactants that are not toxic, carcinogenic, flammable, or environmentally persistent. Using non-hazardous starting materials reduces risks to workers, minimises the cost and complexity of safe handling, and decreases the likelihood of generating harmful by-products during synthesis.

Multi-Step Organic Synthesis

• Retrosynthetic step: A single backwards disconnection in a synthesis plan, in which you identify the functional group of the target and deduce the functional group of the molecule that must have been present one step earlier. Represented by a double-headed arrow (⇒).

Introduction to NMR and TMS Standard

• Chemical shift (δ): The position of an NMR signal on the delta scale, measured in parts per million (ppm) relative to the signal from TMS (assigned δ = 0 ppm). Chemical shift depends on the molecular environment of the nucleus.

Interpreting NMR Spectra and Integration

• Chemical shift (δ / ppm): The position of an NMR peak on the frequency axis, expressed in parts per million relative to TMS at 0 ppm. A higher δ value means the nucleus is more deshielded — it experiences a stronger effective magnetic field and resonates at higher frequency.

Spin-Spin Splitting and Analytical Techniques

• Non-equivalent protons: Protons that are in chemically different environments and therefore resonate at different chemical shifts. Only non-equivalent protons on adjacent carbons cause spin-spin splitting of each other's signals.

Chromatography Types: TLC, CC & GC (Part 2/4)

• Adsorption: Attachment of particles to a surface.
• Partition: Distribution of a substance between two phases according to how soluble it is in each one.

Chromatography Types: TLC, CC & GC (Part 3/4)

• Rf value: Rf = distance travelled by component ÷ distance travelled by solvent front. It is a dimensionless ratio between 0 and 1, characteristic of a compound under a specific set of conditions (same stationary phase, same mobile phase, same temperature).
• Retention time: The time elapsed between injection of the sample and detection of a component at the end of the GC column. Each compound has a characteristic retention time under fixed conditions (same column, same carrier gas flow rate, same oven temperature programme). Used to identify components by comparison with known standards.

Calculations in Chromatography: Rf and Retention Times (Part 1/4)

• Rf value (retardation factor): The ratio of the distance travelled by a substance to the distance travelled by the solvent front, measured from the same baseline. Rf has no units and always lies between 0 and 1 (inclusive). Formula: Rf = distance travelled by substance (mm or cm) ÷ distance travelled by solvent front (mm or cm).