GCSE Chemistry flashcards

About this deck

353 free AQA GCSE Chemistry flashcards across 11 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.

Atomic structure and the periodic table

39 cards from 19 lessons. Each term links to the lesson it comes from.

Electrical Charges in Atoms

• Proton: A positively charged particle found in the nucleus of an atom. Relative charge = +1.
• Electron: A negatively charged particle that orbits the nucleus of an atom. Relative charge = -1.
• Neutron: A particle with no electrical charge found in the nucleus of an atom. Relative charge = 0.
• Atomic number: The number of protons in an atom's nucleus. This defines what element the atom is.

Early atomic models: from spheres to plum pudding

• Sphere model of the atom: The early atomic model proposing that atoms were tiny, solid, indivisible spheres that could not be broken down further.
• Plum pudding model: Thomson's atomic model showing a sphere of positive charge with negative electrons embedded randomly throughout, like raisins in a pudding.

The nuclear model: Rutherford's breakthrough

• Nuclear model: A model of the atom where most of the mass and all positive charge is concentrated in a tiny central nucleus, with electrons orbiting around it

Refining the atom: Bohr, protons and neutrons

• Bohr model: An atomic model where electrons orbit the nucleus at specific, fixed distances called energy levels or electron shells

Understanding Relative Atomic Mass

• Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons. They have identical chemical properties but different masses.
• Relative atomic mass: The weighted mean mass of an atom of an element compared with one twelfth of the mass of a carbon-12 atom.

Electronic Structure and Energy Levels

• Energy levels: Fixed regions around the nucleus where electrons can exist. Also called electron shells. Numbered 1, 2, 3 from innermost to outermost.
• Electron Shell Arrangement: The distribution of electrons in energy levels, written as shell populations such as 2,8,1.

What are mixtures and their properties

• Mixture: A mixture consists of two or more elements or compounds not chemically combined together. The chemical properties of each substance in the mixture are unchanged.
• Chemical properties: The characteristics that determine how a substance reacts with other substances, such as flammability, reactivity with acids, or ability to support combustion.
• Physical separation: Methods that separate mixtures without changing the chemical identity of the substances, exploiting differences in physical properties like magnetism, solubility, or boiling point.

Separation techniques for mixtures

• Filtration: A separation technique that uses filter paper or a membrane to separate insoluble particles from a liquid based on particle size differences.
• Crystallisation: The process of forming solid crystals from a solution by evaporating the solvent, used to separate dissolved solids from liquids.
• Simple distillation: A separation technique that uses differences in boiling points to separate a liquid from a mixture by evaporation and condensation.
• Fractional distillation: An enhanced distillation method using a fractionating column to separate liquids with similar boiling points through repeated evaporation and condensation cycles.
• Chromatography: A separation technique that separates mixtures based on how different substances move through a stationary phase when carried by a mobile phase.

Introduction to Group 1 Elements - The Alkali Metals

• Alkali metals: Elements in Group 1 of the periodic table that have one electron in their outer shell and form alkaline solutions when they react with water. They are highly reactive metals that readily lose their outer electron to form positive ions.
• Outer shell electron: The single electron in the highest energy level of Group 1 atoms. This electron is weakly held because it's far from the nucleus and shielded by inner electrons, making it easy to lose in chemical reactions.
• Key GCSE Chemistry Term: A precise term from Introduction to Group 1 Elements - The Alkali Metals - Group 1 - The periodic table that should be used consistently in exam answers.

Reactions of Group 1 Elements

• Reactivity trend: In Group 1, reactivity increases going down the group because the outer electron is further from the nucleus and easier to remove.

Predicting Group 1 Trends

• Group 1 Elements (Alkali Metals): Elements in the first column of the periodic table with one electron in their outer shell. Include lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr).

Understanding the Periodic Table Structure

• Electron shells: Energy levels around the nucleus where electrons are found. The first shell holds up to 2 electrons, the second and third shells hold up to 8 electrons each.
• Group: A vertical column in the periodic table. Elements in the same group have the same number of electrons in their outer shell, giving them similar chemical properties.
• Period: A horizontal row in the periodic table. Elements in the same period have the same number of electron shells.

Introduction to Group 7 Halogens

• Halogens: Elements in Group 7 of the periodic table that have seven electrons in their outer shell and exist as diatomic molecules (pairs of atoms).
• Halide: An ionic compound formed when a halogen reacts with a metal, where the halogen exists as a negative ion (like Cl^-, Br^-, I^-).
• Displacement reaction: A reaction where a more reactive halogen displaces a less reactive halogen from its compound, based on the reactivity series.

Trends in Group 7 Properties

• Relative molecular mass: The mass of a molecule compared to 1/12 of the mass of a carbon-12 atom. For diatomic molecules like halogens, it's twice the relative atomic mass.

Noble Gases: Structure and Unreactivity

• Noble gases: Elements in Group 0 of the periodic table that are unreactive and do not easily form molecules because their atoms have stable arrangements of electrons.

Introduction to Transition Metal Properties

• Ions with different charges: The ability of transition metals to form ions with different charges by losing different numbers of electrons from their outer shells
• Coloured compounds: Compounds that absorb certain wavelengths of visible light and reflect others, giving them distinctive colours
• Catalyst: A substance that increases the rate of a chemical reaction without being permanently changed or used up in the process

Transition Metal Examples: From Chromium to Copper

• Transition metals: Transition metals are metals in the central part of the periodic table. At GCSE, the key examples are chromium, manganese, iron, cobalt, nickel, and copper.

Properties of Transition Metals vs Group 1 Elements

• Group 1 metals: Alkali metals in the first column of the periodic table that have one electron in their outer shell, making them highly reactive with low melting points and densities.

Examples of Transition Metal Properties

• Catalytic activity: The ability to speed up chemical reactions without being consumed in the process. Many transition metals and their compounds are useful catalysts.

Bonding, structure, and the properties of matter

39 cards from 22 lessons. Each term links to the lesson it comes from.

Understanding Ionic Structure and Bonding

• Ion: A charged particle formed when an atom gains or loses electrons to achieve a stable electron configuration. Metal atoms lose electrons to form positive ions (cations), while non-metal atoms gain electrons to form negative ions (anions).
• Ionic bonding: The electrostatic attraction between oppositely charged ions formed when electrons are transferred from metal atoms to non-metal atoms. The metal loses electrons to form positive ions, while the non-metal gains electrons to form negative ions.
• Giant ionic lattice: A regular three-dimensional arrangement of ions held together by strong electrostatic forces acting in all directions. Each positive ion is surrounded by negative ions, and each negative ion is surrounded by positive ions.

Representing Ionic Structures and Finding Empirical Formulas

• Dot and cross diagram: A 2D representation showing the transfer of electrons from metal atoms to non-metal atoms, using dots and crosses to track electron origins
• Empirical formula: The simplest whole number ratio of atoms or ions in a compound, showing the relative proportions of each element present

Understanding covalent bonds and drawing molecules

• Covalent bond: A chemical bond formed when two atoms share a pair of electrons. The shared electrons are attracted to both nuclei, holding the atoms together strongly.

Types of covalent structures and their representations

• Covalent bond representation: A line drawn between two atoms represents a single covalent bond (one shared pair of electrons). Double bonds use two lines, triple bonds use three lines.
• Polymer representation: Polymers are shown as (monomer unit)ₙ where n is a large number representing how many monomer units are joined together.
• Giant covalent structure: A three-dimensional network of atoms joined by covalent bonds throughout the entire structure, with no discrete molecules present.

Working with molecular diagrams and formulas

• Ball and Stick Model: Represents atoms as spheres (balls) connected by rods (sticks) to show the three-dimensional arrangement of atoms and bonds.

Understanding the Three Types of Chemical Bonds

• Covalent bonding: Chemical bonding between atoms (usually non-metals) that share pairs of electrons in their outer shells. The shared electrons are attracted to both nuclei, creating the bond.
• Metallic bonding: Chemical bonding between metal atoms where electrons are delocalised - they form a 'sea of electrons' that can move freely throughout the metal structure. The positive metal ions are held in place by attraction to this electron sea.

Where Different Types of Bonding Occur

• Alloy: A mixture of metals, or a metal mixed with other elements, that maintains metallic bonding throughout. Examples include steel (iron + carbon) and brass (copper + zinc).

How ionic bonds form through electron transfer

• Electron transfer: The movement of electrons from metal atoms to non-metal atoms during chemical bonding. This creates oppositely charged ions that attract each other.
• Ion charge: The electrical charge on an ion, determined by the difference between the number of protons and electrons. Positive ions have lost electrons, negative ions have gained electrons.
• Noble gas structure: A stable electron arrangement with a full outer shell, like the Group 0 elements. Ions formed by Groups 1, 2, 6, and 7 all achieve noble gas structures.

Physical Properties of Small Molecules

• Small molecules: Substances made of covalently bonded atoms that exist as discrete molecules with weak forces between them. Examples include water, methane, and carbon dioxide.
• Intermolecular forces: Weak forces of attraction between molecules that hold them together in liquids and solids. Much weaker than covalent bonds.

How Molecule Size Affects Properties and Electrical Conductivity

• Electrical conductivity: The ability of a substance to allow electric current to flow through it. Requires mobile charged particles like ions or electrons.

Structure and Properties of Polymers

• Polymer: A substance with very large molecules made from many repeating units (monomers) joined by strong covalent bonds. Examples include polyethene, nylon, and starch.
• Intramolecular forces: Strong covalent bonds that hold atoms together within a polymer molecule. These bonds are very difficult to break.

Giant Metallic Structure and Bonding

• Giant metallic structure: A three-dimensional lattice of metal atoms held together by metallic bonding, containing millions of atoms arranged in a regular pattern.

Layered Structures and Alloys

• Layered structure: The arrangement of metal atoms in regular, repeating layers where atoms are packed closely together in an orderly pattern.

States of Matter and Particle Theory

• Particle Theory: A model that explains the properties of matter by describing substances as made of tiny particles in constant motion. The particles are represented as small solid spheres.
• Melting Point: The temperature at which a solid changes to a liquid (melting) or a liquid changes to a solid (freezing). Energy is needed to break particle arrangements.
• Boiling Point: The temperature at which a liquid changes to a gas (boiling) or a gas changes to a liquid (condensing). Energy breaks all remaining particle attractions.

Energy Changes and Forces Between Particles

• Van der Waals forces: Weak forces between molecules. Creates low melting and boiling points in simple molecular substances like oxygen or methane.

Structure and Properties of Graphite

• Graphite structure: Layers of carbon atoms arranged in hexagonal rings, where each carbon forms three covalent bonds within its layer, with weak forces between layers.
• Delocalised electrons: Electrons that are not bound to specific atoms but can move freely through a structure, creating a 'sea' of mobile charge carriers.

The Structure and Properties of Diamond

• Hardness: The resistance of a material to being scratched, cut, or deformed. Diamond rates 10 on the Mohs hardness scale - the maximum value.

Discovering Graphene and Fullerenes

• Allotrope: Different structural forms of the same element, where atoms are arranged in different ways to give different properties.
• Graphene: A single layer of graphite consisting of carbon atoms arranged in a hexagonal lattice, with each carbon bonded to three others by covalent bonds.
• Fullerenes: Hollow carbon structures where carbon atoms are arranged in rings of five or six atoms, forming cage-like molecules such as buckyballs.

Fullerenes and Carbon Nanotubes

• Buckminsterfullerene (C₆₀): The first fullerene discovered, with 60 carbon atoms arranged in a spherical cage structure resembling a football.
• Carbon Nanotubes: Cylindrical fullerenes with very high length to diametre ratios. They can be thought of as rolled-up sheets of graphite.

Applications of Nanoparticles in Everyday Life

• Nanoparticles: Particles with at least one dimension between 1-100 nanometres that exhibit properties different from bulk materials due to their extremely high surface area to volume ratio.

Particle Size Categories and Scale

• Fine particles (PM2.5): Particles with diametres between 100-2500 nm (1 × 10⁻⁷ m to 2.5 × 10⁻⁶ m). The '2.5' refers to 2.5 micrometres - the upper size limit.
• Coarse particles (PM10): Particles with diametres between 2.5-10 micrometres (2.5 × 10⁻⁶ m to 1 × 10⁻⁵ m). These are often called dust and include pollen, mould spores, and larger pollution particles.

Surface Area to Volume Ratio and Nanoparticle Properties

• Surface Area to Volume Ratio: The ratio of the total surface area of an object to its total volume, showing how much surface is available per unit of volume. Calculated as SA:V or SA/V.

Quantitative chemistry

27 cards from 18 lessons. Each term links to the lesson it comes from.

Calculating Relative Formula Mass

• Relative Atomic Mass (Ar): The average mass of an atom of an element compared to 1/12 of the mass of a carbon-12 atom. Found on the periodic table for each element.
• Relative Formula Mass (Mr): The sum of the relative atomic masses of all atoms in a chemical formula. Calculated by multiplying each element's atomic mass by how many atoms of that element are present, then adding them all up.

Percentage by Mass in Compounds

• Percentage by mass: The mass of a specific element in a compound divided by the total mass of the compound, multiplied by 100 to give a percentage.

Explaining Mass Changes in Open Systems

• Open System: A chemical reaction system where reactants or products can enter or leave the reaction vessel, typically involving gases that escape to the atmosphere.

The Law of Conservation of Mass

• Law of Conservation of Mass: No atoms are lost or made during a chemical reaction, so the mass of the products equals the mass of the reactants.
• Closed System: A reaction container where no substances can enter or leave during the reaction.

Balancing Chemical Equations

• Conservation of mass: The principle that atoms cannot be created or destroyed in chemical reactions, only rearranged into different compounds.
• Subscript: The small number written after an element symbol (like the 2 in H₂O) showing how many atoms of that element are in one molecule or formula unit.
• Coefficient: The normal-sized number written in front of a chemical formula (like the 2 in 2H₂O) showing how many molecules or formula units you have.

Understanding Uncertainty in Chemical Measurements

• Uncertainty: The doubt that exists about the result of any measurement, arising from limitations in measuring instruments and experimental conditions.
• Range: The difference between the highest and lowest values in a set of measurements, used as a measure of uncertainty.

Converting Masses to Moles for Equation Balancing

• Balanced Equation: A chemical equation with the same number of each type of atom on both sides.

Understanding Solution Concentration

• Solute: The substance that dissolves in a solvent to form a solution. This is what you're adding to the liquid.
• Solvent: The liquid that dissolves the solute, usually water in chemistry. This is what does the dissolving.
• Solution: A homogeneous mixture formed when a solute dissolves in a solvent. The final mixture has the same composition throughout.
• Concentration: The amount of solute dissolved in a given volume of solution, commonly measured in grams per cubic decimetre (g/dm³).

Understanding Limiting Reactants

• Limiting reactant: The reactant that is completely used up in a chemical reaction, determining the maximum amount of product that can be formed. Measured by comparing moles of each reactant to the balanced equation.

Understanding Chemical Equations in Terms of Moles

• Mole ratio: The ratio of moles of one substance to moles of another substance in a balanced chemical equation, determined by the coefficients.

Calculating Masses from Balanced Chemical Equations

• Mole: A unit of amount containing 6.022 × 10²³ particles (atoms, molecules, or ions). One mole of any substance has a mass equal to its relative formula mass in grams.

Introduction to Moles and Avogadro's Constant

• Avogadro's Constant: The number of particles in one mole of any substance. Value: 6.02 × 10²³ per mole. Symbol: Nₐ.

Understanding Atom Economy

• Atom Economy: The percentage of starting materials that end up as useful products in a chemical reaction. Calculated using relative formula masses from the balanced equation.

Understanding Percentage Yield and Why Reactions Don't Give 100% Yield

• Percentage Yield: The actual amount of product obtained compared to the theoretical maximum, expressed as a percentage. Calculated using: Percentage Yield = (Actual Yield ÷ Theoretical Yield) × 100%

Understanding Solution Concentration in mol/dm³

• Concentration (mol/dm³): The number of moles of solute dissolved in one cubic decimetre (1 litre) of solution.

Finding Unknown Concentrations Using Reaction Volumes

• Molar ratio: The ratio of moles of reactants and products in a balanced chemical equation

Gas volume fundamentals and molar volume

• Equal moles of gases occupy equal volumes: Equal amounts in moles of gases occupy the same volume under the same conditions of temperature and pressure.
• Molar volume: The volume occupied by one mole of any gas at room temperature and pressure. Always 24 dm³.

Gas volume calculations from balanced equations

• Molar gas volume: The volume occupied by one mole of any gas at room temperature and pressure (RTP). At RTP, this equals 24 dm³ (24,000 cm³).

Chemical changes

44 cards from 20 lessons. Each term links to the lesson it comes from.

Metal Reactions and the Reactivity Series

• Metal Reactivity: The tendency of a metal to lose electrons and form positive ions when reacting with other substances. More reactive metals lose electrons more easily.
• Reactivity Series: An arrangement of metals in order of their reactivity, determined by observing their reactions with water and dilute acids. Used to predict and explain metal reactions.
• Oxidation: The loss of electrons by an atom or ion. When metals react, they are oxidised as they lose electrons to form positive metal ions.

How metals are found and extracted

• Metal ore: A rock containing enough metal compound to make extraction economically worthwhile. Examples include iron ore (haematite, Fe2O3) and aluminium ore (bauxite, Al2O3).
• Native metal: A metal found naturally in its pure form rather than as a compound. Only the least reactive metals like gold, silver, and platinum exist as native metals because they don't readily react with oxygen or other elements.
• Reduction: A chemical reaction where oxygen is removed from a compound. In metal extraction, carbon reduces metal oxides by removing oxygen, leaving the pure metal behind. The carbon itself gets oxidised to carbon dioxide.

Evaluating metal extraction processes

• Reduction with carbon: A method where carbon removes oxygen from metal oxides, used for metals less reactive than carbon (zinc, iron, tin, lead, copper).
• Electrolysis extraction: Using electricity to decompose molten metal compounds, used for very reactive metals (potassium, sodium, calcium, magnesium, aluminium).

Writing Ionic Equations for Displacement Reactions

• Spectator ions: Ions that are present in a reaction mixture but do not take part in the actual reaction. They remain unchanged throughout the process.

Metal Oxides and Oxidation Reactions

• Metal oxide: A compound formed when a metal reacts with oxygen. Contains metal atoms bonded to oxygen atoms.

What Salts Are Produced in Neutralisation?

• Salt Formation Rule: In neutralisation, the salt formed contains: (1) negative ions from the acid, and (2) positive ions from the base, alkali or carbonate.
• Hydrochloric acid (HCl): Always produces chloride salts. The Cl⁻ ions from the acid become part of the salt name.
• Nitric acid (HNO₃): Always produces nitrate salts. The NO₃⁻ ions from the acid become part of the salt name.
• Sulfuric acid (H₂SO₄): Always produces sulfate salts. The SO₄²⁻ ions from the acid become part of the salt name.

Neutralisation Reactions and Salt Formulae

• Neutralisation: A chemical reaction between an acid and a base that produces a salt and water, with pH moving towards 7
• Alkali: A soluble base that releases hydroxide ions (OH⁻) when dissolved in water
• Base: A substance that neutralises acids by accepting hydrogen ions, including metal oxides and hydroxides
• Salt: An ionic compound formed when the hydrogen in an acid is replaced by a metal or ammonium ion

Making Soluble Salts from Acids

• Acid + Metal: Produces a salt and hydrogen gas. Only works with metals above hydrogen in the reactivity series.
• Acid + Metal Oxide: Produces a salt and water only. This is a neutralisation reaction.
• Acid + Metal Hydroxide: Produces a salt and water. Another neutralisation reaction, often called acid-alkali neutralisation.
• Acid + Metal Carbonate: Produces a salt, water and carbon dioxide gas. You'll see fizzing as CO₂ escapes.

Acid-Metal Reactions: Products and Patterns

• Acid-metal reaction: A chemical reaction between an acid and a metal that produces a salt and hydrogen gas. Only occurs with metals above hydrogen in the reactivity series.
• Hydrogen test: A test using a lit splint to identify hydrogen gas. Hydrogen burns with a distinctive 'squeaky pop' sound.

Redox in Acid-Metal Reactions

• Redox reaction: A chemical reaction involving the transfer of electrons, where one species is oxidised (loses electrons) and another is reduced (gains electrons).

Understanding Acid Strength and Concentration

• Strong acid: An acid that completely ionises in aqueous solution, meaning every acid molecule releases its hydrogen ion.
• Weak acid: An acid that only partially ionises in aqueous solution, meaning only some acid molecules release their hydrogen ions.
• Concentrated acid: An acid solution with a large amount of acid dissolved per unit volume of solution.
• Dilute acid: An acid solution with a small amount of acid dissolved per unit volume of solution.

pH Scale and Hydrogen Ion Concentration

• pH Scale: A logarithmic scale from 0 to 14 that measures hydrogen ion concentration, where pH 7 is neutral, below 7 is acidic, and above 7 is basic.
• Neutrality: The condition where hydrogen ion concentration equals hydroxide ion concentration, always occurring at pH 7 in aqueous solutions at 25°C.

Understanding Electrode Reactions in Electrolysis

• Cathode: The negative electrode in electrolysis where positively charged ions gain electrons and are reduced.
• Anode: The positive electrode in electrolysis where negatively charged ions lose electrons and are oxidised.

Writing Half Equations for Electrode Reactions

• Half equation: A chemical equation showing the reaction at one electrode during electrolysis, including electrons (e⁻) as reactants or products
• Reduction half equation: Shows a substance gaining electrons at the cathode, written as: substance + electrons → product
• Oxidation half equation: Shows a substance losing electrons at the anode, written as: substance → product + electrons

Predicting Products in Aqueous Electrolysis

• Reactivity series order: K > Na > Li > Ca > Mg > Al > C > Zn > Fe > Sn > Pb > H > Cu > Ag > Au > Pt Ag > Au > Pt (most reactive to least reactive)

Products at the Cathode and Anode

• Cathode Rule: If the metal is more reactive than hydrogen, hydrogen gas is produced at the cathode. If the metal is less reactive than hydrogen, the metal itself is produced.
• Anode Rule: Oxygen gas is produced at the anode, unless the solution contains halide ions (Cl⁻, Br⁻, I⁻), in which case the halogen gas is produced instead.

Electrolytic extraction of reactive metals

• Electrolytic extraction: The process of extracting metals from their compounds using electrical energy to break ionic bonds. Used when metals are too reactive for carbon reduction or when the metal reacts with carbon to form carbides.
• Cryolite: Sodium aluminium fluoride (Na3AlF6) used as a solvent for aluminium oxide in the Hall-Héroult process. Lowers the melting point and increases conductivity.

Aluminium production and electrode replacement

• Electrolysis: The decomposition of an ionic compound using electrical energy. The compound must be molten or dissolved in water to allow ions to move freely.

Products of Electrolysis in Molten Ionic Compounds

• Ion migration: The movement of ions towards electrodes of opposite charge during electrolysis. Positive ions move to the negative electrode, negative ions move to the positive electrode.

Understanding Electrolytes and Ion Movement

• Electrolyte: A liquid or solution that conducts electricity because it contains free-moving ions. Formed when ionic compounds are melted or dissolved in water.

Energy changes

16 cards from 9 lessons. Each term links to the lesson it comes from.

Understanding Exothermic and Endothermic Reactions

• Exothermic reaction: A chemical reaction that transfers energy to the surroundings, so the temperature of the surroundings increases.
• Endothermic reaction: A chemical reaction that takes in energy from the surroundings, so the temperature of the surroundings decreases.

Examples and Applications of Energy Transfer Reactions

• Combustion: The reaction of a fuel with oxygen that releases energy as heat and light. Complete combustion produces carbon dioxide and water.

Breaking and Forming Bonds

• Bond breaking: The process that requires energy input to overcome the attractive forces holding atoms together in molecules
• Bond forming: The process that releases energy when atoms come together to form new chemical bonds
• Bond energy: The energy required to break one mole of a particular bond in gaseous molecules, measured in kJ/mol

Calculating Energy Changes Using Bond Energies

• Activation energy: A precise term from Calculating Energy Changes Using Bond Energies - The energy change of reactions (HT only) - Exothermic and endothermic reactions that should be used consistently in exam answers.

Drawing and Understanding Reaction Profiles

• Reaction Profile: Energy level diagram that shows how energy changes during a chemical reaction, displaying reactants, products, activation energy, and overall energy change.

Activation Energy and Collision Theory

• Collision Theory: Chemical reactions occur only when reacting particles collide with each other and with sufficient energy to overcome the activation energy barrier.
• Successful Collision: A collision between particles that has enough energy (equal to or greater than activation energy) and correct orientation to cause a reaction.

Introduction to Fuel Cells

• Fuel Cell: An electrochemical device that converts chemical energy from fuel and oxygen directly into electrical energy through oxidation reactions. The fuel is supplied externally and continuously consumed.

How Cells and Batteries Work

• Chemical cell: A device that converts chemical energy into electrical energy through redox reactions between electrodes and an electrolyte. The chemical reactions create a potential difference that drives current through an external circuit.
• Electrode: A conductor through which electricity enters or leaves an electrolyte during chemical reactions. In a cell, one electrode loses electrons (anode) whilst the other gains electrons (cathode).
• Battery: Two or more chemical cells can be combined to make a battery.

Types of Cells and Batteries

• Non-rechargeable cell: A cell where chemical reactions stop permanently when one reactant is used up. The reactions cannot be reversed, so the cell cannot be recharged.
• Rechargeable cell: A cell where chemical reactions can be reversed by supplying external electrical current. This restores the original reactants, allowing the cell to be used again.

The rate and extent of chemical change

14 cards from 10 lessons. Each term links to the lesson it comes from.

What are catalysts and how do they work?

• Enzyme: A biological catalyst made of protein that speeds up specific reactions in living organisms. Each enzyme catalyses only one type of reaction.

What affects how fast reactions happen?

• Rate of reaction: How fast reactants are used up or products are formed in a chemical reaction, measured in g/s or mol/s

Surface Area Effects and Using Proportionality in Collision Theory

• Direct Proportionality: A relationship where one quantity increases by the same factor as another. If A doubles, B doubles. Written as A ∝ B.

Calculating Mean Rate of Reaction

• Mean Rate of Reaction: The average speed at which a chemical reaction proceeds, calculated by dividing the change in quantity by the time taken. Measured in units like g/s, cm³/s, or mol/s.

Rate Graphs and Tangents

• Rate Graph: A graph showing how the amount of reactant or product changes over time during a chemical reaction
• Tangent Line: A straight line that touches a curve at exactly one point, showing the instantaneous rate of change at that moment

How Temperature Changes Affect Chemical Equilibrium

• Le Chatelier's principle: When a system at equilibrium experiences a change in conditions, the equilibrium position shifts to counteract that change.

How concentration changes affect equilibrium

• Concentration disturbance: Any change to the amount of reactant or product that disrupts the equilibrium state, forcing the system to shift position
• Rule 1: Increase reactant concentration: The equilibrium shifts towards the products (right) to consume the excess reactant until a new equilibrium is established
• Rule 2: Decrease product concentration: The equilibrium shifts towards the products (right) to replace the removed product until a new equilibrium is established

Predicting equilibrium shifts from data

• Equilibrium shift: The change in equilibrium position when conditions are altered, moving to favour either reactants or products
• Dynamic Equilibrium: A state in a closed system where forward and reverse reactions continue at equal rates.

How Pressure Changes Affect Equilibrium Position

• Pressure-Equilibrium Rule: Increased pressure shifts equilibrium towards the side with fewer gas molecules. Decreased pressure shifts equilibrium towards the side with more gas molecules.

Understanding Equilibrium Systems and Le Chatelier's Principle

• Reversible Reaction: A chemical reaction that can proceed in both forward and backward directions under the same conditions. Shown using the equilibrium arrow symbol ⇌.

Organic chemistry

40 cards from 19 lessons. Each term links to the lesson it comes from.

Separating Crude Oil by Fractional Distillation

• Evaporation: The change of state from liquid to gas when molecules gain enough energy to escape the liquid surface.
• Condensation: The change of state from gas to liquid when molecules lose energy and come together to form a liquid.

From Crude Oil to Modern Life - Fuels and Petrochemicals

• Petrochemicals: Chemical compounds derived from petroleum that serve as feedstock for manufacturing synthetic materials, including polymers, solvents, lubricants, and detergents.

What is crude oil and where does it come from?

• Crude oil: A finite fossil fuel found in rocks, formed from the remains of ancient biomass (mainly plankton) buried in mud millions of years ago
• Biomass: Living or recently dead organic material, in crude oil's case mainly ancient plankton that died millions of years ago
• Hydrocarbon: A compound containing only hydrogen and carbon atoms, which makes up most of the compounds in crude oil

Understanding alkanes - the main hydrocarbons in crude oil

• Alkanes: Saturated hydrocarbons that contain only single bonds between carbon atoms, with the general formula CnH2n+2

Representing and recognising alkane molecules

• Alkane: Saturated hydrocarbons containing only single bonds between carbon atoms, with the general formula CnH2n+2. They are the simplest organic compounds.

Introduction to Cracking and Its Methods

• Cracking: The process of breaking down large hydrocarbon molecules into smaller, more useful ones using heat and/or catalysts
• Catalytic Cracking: A cracking method using moderate temperature (450°C) and a zeolite catalyst to break C-C bonds in hydrocarbons
• Steam Cracking: A cracking method using very high temperature (800°C) and steam to break down hydrocarbons without a catalyst
• Alkenes: Unsaturated hydrocarbons containing C=C double bonds, produced as one of the main products of cracking

Uses and Importance of Cracking Products

• Cracking Products: Smaller hydrocarbons including alkanes (for fuels) and alkenes (for chemical feedstock) produced by breaking down larger molecules
• Polymerisation: The process where many small alkene molecules (monomers) join together to form long-chain polymers

Introduction to Alcohols and Their Structure

• Alcohol: Organic compounds containing the functional group -OH (hydroxyl group). This -OH group determines their chemical behaviour and properties.
• Homologous Series: A family of organic compounds with the same functional group and similar chemical properties. Each member differs from the next by a CH₂ unit.

Fermentation and Alcohol Production

• Fermentation: An anaerobic process where yeast converts sugar into ethanol and carbon dioxide

Introduction to Alkenes and Their General Formula

• Alkene: An unsaturated hydrocarbon containing at least one carbon-carbon double bond (C=C). Members of the alkene homologous series have the general formula CnH2n.
• Unsaturated: A molecule that contains fewer hydrogen atoms than the corresponding saturated compound with the same number of carbon atoms. Contains at least one C=C double bond.

Introduction to Carboxylic Acids

• Carboxylic acid: Organic compounds that contain the functional group -COOH, which gives them acidic properties and the ability to form salts with metals and bases.
• Methanoic acid: The simplest carboxylic acid with formula HCOOH, found naturally in ant stings and nettle stings.
• Ethanoic acid: A carboxylic acid with formula CH₃COOH, commonly known as the acid in vinegar.
• Propanoic acid: A carboxylic acid with formula CH₃CH₂COOH, used as a food preservative.
• Butanoic acid: A carboxylic acid with formula CH₃CH₂CH₂COOH, responsible for the smell of rancid butter.

Chemical Properties of Carboxylic Acids

• Ester: Organic compounds formed when carboxylic acids react with alcohols, containing the functional group -COO- and often having fruity smells.

Introduction to Alkenes and Their Structure

• Functional Group: A specific arrangement of atoms that determines the chemical reactions of organic compounds. In alkenes, this is the C=C double bond.
• Incomplete Combustion: Burning in limited oxygen supply, producing carbon monoxide (CO) and carbon particles instead of just carbon dioxide (CO₂).

What are Addition Polymers?

• Addition Polymer: A large molecule formed when many small molecules (monomers) containing C=C bonds join together in a chain reaction, with no other products formed.
• Monomer: A small molecule containing a C=C double bond that can join with many others to form a polymer chain.

Understanding Condensation Polymerisation

• Condensation polymerisation: A polymerisation reaction where monomers with two functional groups join together, eliminating small molecules such as water, and the reactions are called condensation reactions.
• Elimination: The removal of a small molecule (usually water) when two functional groups react together to form a new bond.

Formation of Polyesters from Ethanediol and Hexanedioic Acid

• Ethanediol: A diol monomer with the formula HOCH₂CH₂OH. It contains two alcohol (-OH) functional groups, one at each end of the molecule.
• Hexanedioic acid: A dicarboxylic acid monomer with the formula HOOC(CH₂)₄COOH. It contains two carboxylic acid (-COOH) functional groups, one at each end of the molecule.

Amino Acids and Polypeptide Formation

• Amino acid: Organic molecules containing both an amino group (-NH₂) and a carboxyl group (-COOH) in the same molecule. The amino group is basic while the carboxyl group is acidic.
• Polypeptide: A polymer formed from many amino acids joined by peptide bonds through condensation polymerisation. Contains repeating -CO-NH- links along the backbone.

DNA Structure and Function

• DNA (Deoxyribonucleic Acid): A large molecule essential for life that encodes genetic instructions for the development and functioning of living organisms and viruses
• Nucleotide: The monomer units that make up DNA polymer chains, with four different types forming the genetic code

Naturally Occurring Polymers and Their Monomers

• Nucleotides: The monomers of DNA and RNA, each containing a phosphate group, a sugar molecule, and a nitrogen-containing base. DNA uses A, T, G and C, while RNA uses A, U, G and C.
• Amino acids: The monomers of proteins, containing an amino group and a carboxyl group, with 20 different types creating protein diversity
• Glucose: The monomer of starch and cellulose, a simple sugar molecule with the formula C₆H₁₂O₆

Chemical analysis

27 cards from 17 lessons. Each term links to the lesson it comes from.

Paper Chromatography - Principles and Applications

• Stationary Phase: The paper in chromatography that remains fixed. Substances interact with it through attraction forces, causing different substances to move at different rates.
• Mobile Phase: The solvent that moves up the paper, carrying dissolved substances with it. Different substances have different solubilities, so they travel at different speeds.

Chromatogram Analysis and Rf Values

• Chromatogram: The pattern of separated substances produced by chromatography, showing spots at different distances from the origin.
• Rf value: The ratio of distance moved by a substance to distance moved by the solvent front in chromatography. Always between 0 and 1.
• Solvent front: The furthest point reached by the solvent as it moves up the chromatography paper.

Understanding Formulations

• Formulation: A mixture that has been designed as a useful product, where each component has a specific purpose and is present in carefully measured quantities.

What is a Pure Substance?

• Pure substance (chemistry): A single element or compound, not mixed with any other substance
• Pure substance (everyday): A substance that has had nothing added to it, so it is unadulterated and in its natural state

Using Melting and Boiling Points to Test Purity

• Pure substance: A substance containing only one type of element or compound with no other substances mixed in. Pure substances have fixed melting and boiling points that act like fingerprints for identification.
• Melting point range: The temperature range over which an impure substance melts, starting when melting begins and ending when it's completely liquid. Pure substances melt at a single, sharp temperature.

Testing for Oxygen Gas

• Glowing splint test: A qualitative test for oxygen gas where a glowing wooden splint is inserted into the test gas. If oxygen is present, the splint relights with a flame.

Testing for hydrogen gas

• Hydrogen test procedure: Hold a burning splint at the open end of a test tube containing the gas. Hydrogen burns rapidly with a distinctive 'pop' sound.

Testing for Carbon Dioxide Gas

• Limewater: An aqueous solution of calcium hydroxide, Ca(OH)₂(aq), used as a chemical test for carbon dioxide gas. It appears as a clear, colourless solution.

Testing for Chlorine Gas

• Chlorine gas test: A chemical test using damp litmus paper that turns white when exposed to chlorine gas due to bleaching
• Bleaching: The removal of colour from a substance by breaking down the molecules responsible for colour
• Chlorine test: A gas test in which damp litmus paper is exposed to chlorine. A positive result is that the paper is bleached and turns white.

Testing for Halide Ions with Silver Nitrate

• Halide ions: Negative ions formed when Group 7 elements (halogens) gain one electron to achieve a stable electron configuration. The three you need to know are chloride (Cl⁻), bromide (Br⁻), and iodide (I⁻).
• Precipitate colours: Silver chloride = white, silver bromide = cream, silver iodide = yellow. Remember: the colours get darker as you go down Group 7.

Testing for Metal Ions with Sodium Hydroxide

• Sodium hydroxide test: A chemical test that identifies metal ions by the colour and behaviour of precipitates formed when NaOH(aq) is added

Coloured Precipitates and Chemical Equations

• Precipitate: An insoluble solid that forms when two solutions are mixed together. The solid appears to 'fall out' of solution.

Testing for Carbonates with Acids

• Carbonate test: A chemical test where dilute acid is added to a suspected carbonate. If carbon dioxide gas is produced (detected by limewater), carbonates are present.

Testing for Sulfate Ions

• Sulfate test procedure: Add a few drops of dilute hydrochloric acid to the test solution, then add barium chloride solution. A white precipitate confirms sulphate ions are present.

Flame Tests for Metal Ion Identification

• Flame test: A qualitative analytical technique that identifies metal ions by the characteristic colour they produce when heated in a flame. Used to identify Group 1 and Group 2 metals and transition metals like copper.

Introduction to Flame Emission Spectroscopy

• Flame emission spectroscopy: An instrumental method that identifies metal ions by analysing the characteristic light they emit when heated in a flame. Used when you need to identify which metals are present in a solution and measure their concentrations.
• Line spectrum: A pattern of coloured lines produced when light from excited atoms is passed through a spectroscope. Each element produces its own unique pattern, like a fingerprint made of light.

Interpreting Flame Emission Spectroscopy Results

• Emission Wavelength: The specific wavelength of light emitted by an excited metal ion, measured in nanometres (nm). Each metal ion has characteristic wavelengths that act like a fingerprint for identification.
• Reference Data: Known emission wavelengths for different metal ions, provided as tables or charts, used to identify unknown samples by matching peak positions.

Chemistry of the atmosphere

40 cards from 18 lessons. Each term links to the lesson it comes from.

How Earth's Early Atmosphere Formed

• Volcanic outgassing: The process by which gases trapped in molten rock are released during volcanic eruptions, contributing to atmospheric formation.
• Atmospheric Pollutant: A substance released into air that can cause environmental or health harm.

Changes in Early Atmospheric Composition

• Ocean formation: The process by which water vapour in the early atmosphere condensed to form liquid water as Earth cooled, creating the first seas and oceans around 3.8 billion years ago.
• Carbonate precipitation: The formation of solid carbonate compounds when dissolved CO2 reacts with minerals in seawater, creating sediments that remove carbon from the atmosphere permanently.

How Photosynthesis Increased Atmospheric Oxygen

• Photosynthesis: The process by which plants and algae convert carbon dioxide and water into glucose and oxygen using light energy. The oxygen is released as a waste product.

How Earth's Atmosphere Changed Over Time

• Outgassing: The release of gases from molten rock during volcanic eruptions, which formed Earth's early atmosphere
• Ozone layer: A layer of ozone gas in the upper atmosphere that absorbs harmful ultraviolet radiation, protecting life on Earth's surface

Formation of Carbon-Rich Deposits

• Limestone formation: The process where marine organisms extract CO₂ from seawater to build calcium carbonate shells and skeletons. When they die, these accumulate on the sea floor and compress into limestone rock over millions of years.
• Hydrocarbon formation: The process where ancient marine microorganisms are buried under sediments and transformed by heat and pressure into crude oil and natural gas over millions of years.

The composition of today's atmosphere

• Nitrogen: An unreactive gas that makes up approximately four-fifths (80%) of the atmosphere. It exists as diatomic molecules (N2) and is largely inert under normal conditions.
• Oxygen: A reactive gas essential for respiration and combustion that makes up approximately one-fifth (20%) of the atmosphere. It exists as diatomic molecules (O2).
• Carbon dioxide: A trace gas making up about 0.04% of the atmosphere, essential for photosynthesis and a key greenhouse gas.
• Water vapour: Variable amounts of gaseous water in the atmosphere, typically 0-4% depending on location and weather conditions.

The stability of atmospheric composition

• Atmospheric composition: The mixture of gases that make up Earth's atmosphere, consisting of approximately 78% nitrogen, 21% oxygen, 0.04% carbon dioxide, and trace amounts of other gases.
• Negative feedback: A self-regulating mechanism where an increase in one factor triggers processes that reduce that factor, maintaining stability.

Human Activities That Increase Greenhouse Gases

• Greenhouse gases: Gases in the atmosphere that absorb infrared radiation from Earth's surface and re-emit it, trapping heat and warming the planet.

Evaluating Climate Change Evidence and Scientific Communication

• Evidence Quality Criteria: Scientific evidence is evaluated based on methodology (how data was collected), sample size (amount of data), reproducibility (can others get the same results), and peer review status (has it been checked by experts).
• Scientific Uncertainty: The degree of doubt in scientific measurements and predictions. Arises from measurement limitations, natural variation, and incomplete understanding of complex systems.

What are greenhouse gases and why are they important?

• Greenhouse effect: The process by which greenhouse gases in the atmosphere absorb infrared radiation from Earth's surface and re-emit it back down, warming the planet.
• Greenhouse gas: A gas that absorbs infrared radiation emitted from the Earth's surface and re-emits it in all directions, trapping heat in the atmosphere. Key greenhouse gases include water vapour, carbon dioxide and methane.

How does the greenhouse effect work?

• Short Wavelength Radiation: High energy electromagnetic radiation from the Sun, including visible light and ultraviolet. Most greenhouse gases are transparent to this radiation.
• Long Wavelength Radiation: Lower energy electromagnetic radiation, mainly infrared, emitted by warm objects like Earth's surface. Greenhouse gases absorb this radiation strongly.
• Carbon Dioxide (CO₂): The most important greenhouse gas from human activities. Produced by burning fossil fuels, deforestation, and respiration. Makes up about 0.04% of the atmosphere.
• Methane (CH₄): A greenhouse gas about 25 times more effective than CO₂ at trapping heat. Produced by agriculture, landfills, and natural gas extraction.

Understanding Carbon Footprints

• Carbon footprint: The total amount of greenhouse gases emitted over the full life cycle of a product, service or event.
• CO₂ equivalent: A unit that expresses the warming potential of different greenhouse gases compared to carbon dioxide over a 100-year period.
• Life cycle assessment: The systematic evaluation of environmental impacts associated with all stages of a product's life, from raw material extraction through production, use and disposal.

Understanding Global Climate Change

• Global Climate Change: Long-term changes in Earth's climate patterns caused by increased average global temperatures. Results from enhanced greenhouse effect due to human activities.

Scale and Implications of Climate Change

• Sea Level Rise: The increase in global ocean levels caused by thermal expansion of seawater and melting of land-based ice sheets and glaciers.
• Tipping Points: Critical thresholds in the climate system where small changes trigger large, often irreversible shifts in Earth's systems.

Fuel Composition and Combustion Products

• Particulates: Microscopic solid particles suspended in the atmosphere, including soot, unburned fuel fragments, and other combustion residues.

Atmospheric Pollutants from Fuel Combustion

• Carbon monoxide (CO): A toxic, colourless gas produced during incomplete combustion of carbon-containing fuels when oxygen supply is limited
• Soot (carbon particles): Fine black particles of unburnt carbon produced during incomplete combustion, especially of hydrocarbon fuels
• Sulfur dioxide (SO₂): An acidic gas produced when sulfur impurities in fossil fuels react with oxygen during combustion
• Nitrogen oxides (NOₓ): Gases including nitrogen monoxide (NO) and nitrogen dioxide (NO₂) formed when nitrogen and oxygen in air react at high combustion temperatures

Toxic and Respiratory Pollutants

• Carbon monoxide: A toxic gas with the formula CO that is colourless and odourless, making it extremely dangerous as it cannot be easily detected by humans.
• Sulfur dioxide: A gas with formula SO2 produced when sulfur-containing fuels burn, causing respiratory problems and contributing to acid rain formation.
• Nitrogen oxides: Gases including NO and NO2 formed when nitrogen and oxygen react at high temperatures, particularly in car engines and power stations.

Particulate Matter and Global Dimming

• Particulate Matter (PM): Tiny solid particles and liquid droplets suspended in air, ranging from visible dust to microscopic particles smaller than bacteria. Measured by diametre in micrometres (μm).
• Global Dimming: The reduction in solar radiation reaching Earth's surface due to particulates in the atmosphere reflecting and absorbing sunlight. Measured as a percentage decrease in solar irradiance.

Using resources

63 cards from 24 lessons. Each term links to the lesson it comes from.

What is Potable Water and How Do We Make It?

• Potable water: Water that is safe to drink, containing sufficiently low levels of dissolved salts and microbes. Not chemically pure as it contains dissolved substances.
• Fresh water: Water with low levels of dissolved substances, typically from rain that collects in rivers, lakes and underground sources.
• Sterilisation: The process of killing harmful microbes in water using agents like chlorine, ozone or ultraviolet light.

Water Treatment for Different Sources

• Ground water treatment: The process of removing bacteria, sediment and small amounts of dissolved substances from water found underground
• Salty water treatment: The process of removing large quantities of dissolved salts from seawater or brackish water to make it potable
• Distillation: A desalination method that heats seawater to create steam, then condenses the steam back to pure water, leaving salts behind
• Reverse osmosis: A desalination method that uses pressure to force water molecules through a semi-permeable membrane, blocking salt ions

Why We Need New Methods to Extract Metals

• Sustainable extraction: Methods of obtaining metals that minimise environmental damage and can continue long-term without depleting resources or destroying ecosystems
• Leachate solution: A liquid containing dissolved metal compounds produced when bacteria break down metal ores through chemical processes

Processing and Evaluating Alternative Metal Extraction Methods

• Bioleaching: The use of bacteria to extract metals from ores by converting insoluble metal sulphides into soluble metal sulphates. The bacteria obtain energy by oxidising the sulfur compounds.
• Phytoremediation: The use of plants to absorb and concentrate metals from contaminated soil. The metal-rich plant material is then harvested and processed to extract the metals.

Physical Treatment of Waste Water

• Screening: The first stage of sewage treatment where large objects and debris are removed using metal bars or mesh screens
• Grit removal: The second stage where sand, gravel and other heavy particles are separated from the waste water flow
• Sedimentation: The process where suspended solid particles settle out of sewage under the influence of gravity, separating it into two distinct layers

Biological Treatment of Waste Water

• Anaerobic digestion: The breakdown of organic matter by bacteria in sealed containers without oxygen, producing biogas and digestate
• Aerobic biological treatment: The use of oxygen-requiring bacteria to break down organic pollutants in wastewater, producing clean water and carbon dioxide

Earth's Resources and Human Needs

• Finite resources: Natural materials that exist in limited quantities and cannot be replaced once used up. Examples include fossil fuels (coal, oil, gas) and metal ores (iron, copper, aluminium).
• Renewable resources: Natural materials that can be replenished by natural processes over time. Examples include timber from forests, fresh water from the water cycle, and energy from wind or solar power.

Data Analysis and Sustainable Development

• Order of Magnitude: A factor of ten difference between values. Used to compare quantities that differ by powers of 10.
• Sustainable Development: A precise term from Data Analysis and Sustainable Development - Using the Earth's resources and sustainable development - Using the Earth's resources and obtaining potable water that should be used consistently in exam answers.

The Four Stages of Life Cycle Assessment

• Life Cycle Assessment (LCA): A systematic method for evaluating the environmental impact of a product throughout its entire life cycle, from raw material extraction to final disposal.

Challenges and Limitations of Life Cycle Assessment

• Quantifiable impacts: Environmental effects that can be measured directly using instruments and expressed as numerical values, such as water consumption in litres or energy use in kilojoules.
• Value judgements: Subjective decisions about the relative importance or severity of different environmental impacts, which cannot be determined by scientific measurement alone.
• Selective LCA: An abbreviated Life Cycle Assessment that examines only certain stages or environmental impacts of a product, rather than the complete lifecycle from raw materials to disposal.

The Three Rs: Reduce, Reuse, Recycle

• Raw Material Extraction: The process of obtaining natural resources from Earth through mining (underground) or quarrying (surface). Creates environmental damage including habitat destruction, water pollution, and landscape scarring.
• Reduce: Using fewer resources and materials in the first place. The most effective strategy because it prevents environmental impact rather than managing it after the fact.
• Reuse: Using items again for their original purpose or finding new uses for them. More effective than recycling because it doesn't require reprocessing energy.
• Recycle: Processing waste materials into new products. Less effective than reduce and reuse because it still requires energy for collection, sorting, and reprocessing.

Recycling in Practice: Glass and Metals

• Glass recycling: The process of crushing waste glass and melting it to form new glass products. The crushed glass is called cullet.

Types of Polymers and Their Production

• Thermosoftening polymer: Polymer with weak intermolecular forces between separate chains that melts when heated and can be reshaped repeatedly
• Thermosetting polymer: Polymer with strong covalent cross-links between chains that does not melt when heated but may char or decompose at high temperatures

Glass and Clay Ceramics

• Soda-lime glass: The most common type of glass, made by heating a mixture of sand (silicon dioxide), sodium carbonate, and limestone. Used for windows, bottles, and everyday glassware.
• Borosilicate glass: Heat-resistant glass made from sand and boron trioxide, which melts at higher temperatures than soda-lime glass. Used for laboratory equipment and ovenware.
• Clay ceramics: Materials made by shaping wet clay and then heating in a furnace. Examples include pottery, bricks, and tiles.

Material Properties and Composites

• Metals: Crystalline materials with metallic bonding that conduct electricity and heat well, are malleable and ductile
• Ceramics: Non-metallic, inorganic materials with ionic or covalent bonding, typically hard and brittle with high melting points
• Polymers: Large molecules made of repeating units (monomers) joined by covalent bonds, typically lightweight and flexible
• Matrix: The continuous phase that surrounds the reinforcement, transfers loads between fibres, and protects them from damage
• Reinforcement: The dispersed phase that provides the main strength and stiffness to the composite, usually in fibre or particle form

Common Alloys and Their Properties

• Bronze: An alloy of copper and tin, historically important and still used today for sculptures, ship propellers, and musical instruments.
• Brass: An alloy of copper and zinc, known for its golden appearance and resistance to corrosion, used in musical instruments and decorative items.
• High Carbon Steel: Steel containing 0.6-1.5% carbon, making it strong but brittle, used for cutting tools and springs.
• Low Carbon Steel: Steel containing less than 0.3% carbon, making it softer and more easily shaped, used for car bodies and construction.

Analysing Alloy Compositions and Applications

• Alloy composition: The exact percentage of each element in an alloy, which determines its final properties and suitable applications.
• Property modification: The deliberate change in an alloy's characteristics achieved by adjusting the type and amount of elements added to the base metal.

Understanding Corrosion and Rusting Experiments

• Corrosion: The destruction of materials by chemical reactions with substances in the environment. Measured by mass loss or structural damage over time.
• Rusting: A specific type of corrosion affecting iron and steel when exposed to oxygen and water. Forms hydrated iron(III) oxide with the formula Fe₂O₃·xH₂O.

Preventing Corrosion Through Barrier Methods

• Barrier method: A corrosion prevention technique that creates a physical layer between the metal and corrosive substances like oxygen and water
• Oxide layer: A thin coating of metal oxide that forms naturally on some metals and protects them from further oxidation

Sacrificial Protection and Galvanising

• Sacrificial Protection: A method of preventing corrosion where a more reactive metal is attached to a less reactive metal. The more reactive metal corrodes instead, 'sacrificing' itself to protect the other metal.
• Galvanising: The process of coating iron or steel with a thin layer of zinc to prevent rusting. The zinc provides both a physical barrier and sacrificial protection.

NPK fertilisers and their importance

• Macronutrients: Essential elements that plants need in large quantities for healthy growth - nitrogen, phosphorus and potassium are the three most important
• NPK fertiliser: A formulation of salts containing nitrogen, phosphorus and potassium compounds in appropriate percentages to improve plant growth and agricultural productivity
• Agricultural productivity: The efficiency of food production measured by crop yield per unit area of land

Processing phosphate rock to make fertilisers

• Phosphate rock: A naturally occurring mineral containing calcium phosphate, Ca₃(PO₄)₂, which is insoluble in water and cannot be used directly as fertiliser
• Nitric acid treatment: Produces calcium nitrate and phosphoric acid - provides both nitrogen and phosphorus nutrients
• Sulfuric acid treatment: Produces calcium sulphate and phosphoric acid - provides phosphorus and sulfur nutrients
• Phosphoric acid treatment: Produces calcium dihydrogen phosphate - provides concentrated phosphorus nutrients

Industrial vs laboratory production of fertilisers

• Continuous process: Industrial method where reactants are continuously fed in and products continuously removed, allowing 24/7 operation.
• Batch process: Laboratory method where a fixed amount of reactants is processed in discrete batches, then the apparatus is cleaned and reloaded.

Introduction to the Haber Process

• Nitrogen Source: Nitrogen gas is obtained from the air by fractional distillation of liquid air, taking advantage of nitrogen's lower boiling point compared to oxygen
• Hydrogen Source: Hydrogen gas is produced from natural gas (methane) by steam reforming: CH₄ + H₂O → CO + 3H₂
• Iron Catalyst: A metal catalyst that provides an alternative reaction pathway with lower activation energy, allowing nitrogen and hydrogen to react at a reasonable rate

Key ideas

4 cards from 2 lessons. Each term links to the lesson it comes from.

Atoms, Elements and the Periodic Table

• Atom: The smallest particle of an element that retains its chemical properties, consisting of a nucleus containing protons and neutrons, surrounded by electrons
• Element: A substance made of only one type of atom, which cannot be broken down into simpler substances by chemical means
• Periodic Table: An arrangement of all known elements in order of increasing atomic number, showing repeating patterns in properties

Chemical Reactions and Energy

• Energy Conservation: The principle that energy cannot be created or destroyed in chemical reactions, only transferred between different forms.