BGCSE Chemistry 0570 cheat sheet

All matter is made of particles (atoms, ions, molecules); the spaces between particles are mostly empty and are largest in gases.

Diffusion is the net movement of particles from higher to lower concentration; rate rises with temperature and falls with particle mass (lighter gases diffuse faster).

Brownian motion of smoke or pollen grains is random jostling caused by unbalanced collisions with fast-moving fluid particles, giving evidence that particles move.

Dissolving is not melting: solute particles separate and spread between solvent particles, but no new substance forms.

Diffusion in liquids is slower than in gases because particles are closer together and collide more frequently.

Particle spacing explains compression: gases are easily compressed, liquids and solids only slightly because their particles are already close.

A pure substance has one type of particle; a mixture contains two or more types not chemically joined and separable by physical means.

Melting: solid to liquid; boiling: liquid to gas (throughout the liquid); evaporation: liquid to gas at the surface below boiling point.

Melting and boiling are endothermic; freezing and condensation are exothermic, with no temperature change during a change of state.

Heating a solid: temperature rises until the melting point, then stays constant while latent heat breaks the lattice.

Kinetic theory: particle kinetic energy increases with temperature; average kinetic energy is proportional to absolute temperature in kelvin.

Gas pressure comes from collisions of particles with container walls; raising temperature or reducing volume increases pressure.

Particles in a solid vibrate about fixed positions; in a liquid they slide past each other; in a gas they move rapidly and randomly.

Volatile liquids evaporate faster because more particles have enough energy to escape at the surface.

Proton mass 1 charge +1; neutron mass 1 charge 0; electron mass about 1/1840 charge -1.

Atomic number Z = number of protons; mass number A = protons + neutrons; neutrons = A - Z.

Isotopes have the same Z but different A; they have identical chemical properties but different masses and physical behaviour.

Electrons occupy shells filled in order 2, 8, 8, 18; the outer shell holds at most 8 electrons.

Ions form by loss or gain of electrons: metals lose electrons (cations), non-metals gain electrons (anions).

Relative atomic mass is the weighted mean of isotope masses on the carbon-12 scale.

Electron configuration determines group (outer electrons) and period (number of occupied shells).

Groups are vertical columns with the same number of outer electrons; periods are horizontal rows with the same number of shells.

Group I metals become more reactive down the group; Group VII non-metals become less reactive down the group.

Ionic bonding: metal transfers electrons to non-metal; ions held in a giant lattice; high melting point; conducts when molten or aqueous.

Covalent bonding: shared electron pairs between non-metals; simple molecules have weak intermolecular forces and low melting points.

Metallic bonding: positive ions in a sea of delocalised electrons, giving conductivity and malleability.

Noble gases (Group VIII/0) are unreactive because they have full outer shells.

Giant covalent structures (diamond, graphite, silicon dioxide) have very high melting points; graphite conducts because of delocalised electrons.

Moles n = mass / molar mass; for gases at room temperature and pressure, n = volume / 24 dm3.

Concentration in mol/dm3 = moles / volume in dm3; mass concentration g/dm3 = mass / volume in dm3.

Avogadro constant = 6.02 x 10^23 particles per mole.

Percentage yield = (actual yield / theoretical yield) x 100; atom economy = (molar mass of desired product / total molar mass of reactants) x 100.

Empirical formula from masses or percentages: divide each by its relative atomic mass, then divide by the smallest ratio.

Balance equations by adjusting coefficients only, so atoms of each element and the total charge are equal on both sides.

Water of crystallisation is found by heating to constant mass and comparing mass lost with anhydrous mass.

Acid + base gives salt + water; acid + metal gives salt + hydrogen; acid + carbonate gives salt + water + carbon dioxide.

pH below 7 is acidic, 7 neutral, above 7 alkaline; pH measures hydrogen ion concentration on a logarithmic scale.

Neutralisation: H+ + OH- gives H2O; the salt is named from the metal (or ammonium) cation and the acid anion.

Solubility rules: all nitrates and sodium, potassium and ammonium salts are soluble; most chlorides soluble except silver and lead; most sulfates soluble except barium, calcium and lead.

Rate increases with concentration, pressure (gases), surface area and temperature, and with a catalyst.

Collision theory: reaction requires collisions with energy at least the activation energy and correct orientation.

A catalyst lowers activation energy and is not used up; enzymes are biological catalysts.

Oxidation is loss of electrons; reduction is gain of electrons (OIL RIG); oxidising agents gain electrons.

Electrolysis of molten lead bromide gives lead at the cathode and bromine at the anode.

At the cathode, cations gain electrons (reduction); at the anode, anions lose electrons (oxidation).

In aqueous solution, the least reactive ion is usually discharged: hydrogen is produced at the cathode unless a reactive metal is present.

Electrolysis of concentrated sodium chloride solution gives hydrogen at the cathode, chlorine at the anode and sodium hydroxide in solution.

Electroplating uses the object as the cathode and the plating metal as the anode, with a salt of the plating metal as electrolyte.

Oxidation state rules: elements 0; oxygen usually -2; hydrogen usually +1; sum equals the ion charge.

Exothermic reactions release energy (negative enthalpy change); endothermic reactions absorb energy (positive enthalpy change).

Bond breaking is endothermic; bond making is exothermic; overall sign is the balance of the two.

Activation energy is the minimum energy needed for a reaction; catalysts lower it.

Le Chatelier: raising temperature shifts equilibrium in the endothermic direction; raising pressure shifts toward fewer gas moles.

Increasing the concentration of a reactant shifts equilibrium toward products; a catalyst changes rate but not position of equilibrium.

Haber process: N2 + 3H2 gives 2NH3; exothermic forward, so moderate temperature (about 450 C) and high pressure favour yield.

Enthalpy change can be measured by calorimetry: q = mc(delta T), with m in grams and c = 4.18 J/(g K) for water.

Reactivity series: potassium, sodium, calcium, magnesium, aluminium, zinc, iron, tin, lead, copper, silver, gold (descending).

Metals above carbon are extracted by electrolysis; metals below carbon are reduced with carbon or carbon monoxide.

Iron is extracted in a blast furnace from haematite using coke, limestone and hot air; the product is impure pig iron.

Aluminium is extracted by electrolysis of alumina dissolved in molten cryolite, requiring large amounts of electricity.

Alloys are mixtures of a metal with other elements; they are harder than pure metals because layers cannot slide easily.

Anodising thickens the protective oxide layer on aluminium; galvanising coats iron with zinc to prevent rusting.

Copper is purified by electrolysis using impure copper as the anode and pure copper as the cathode.

Alkanes CnH2n+2 are saturated; alkenes CnH2n contain a carbon-carbon double bond and are unsaturated.

Alkenes decolourise bromine water; alkanes do not, so this is a test for unsaturation.

Complete combustion gives carbon dioxide and water; incomplete combustion gives carbon monoxide and soot.

Fractional distillation of crude oil separates fractions by boiling point; larger molecules have higher boiling points.

Cracking breaks long chains into shorter alkanes and alkenes using heat and a catalyst.

Ethanol is made by fermentation of sugars with yeast or by hydration of ethene with steam and a catalyst.

Polymers form by addition (alkene monomers) or condensation (monomers with two functional groups, releasing water).

Clean dry air is about 78% nitrogen, 21% oxygen, 0.9% argon and 0.04% carbon dioxide by volume.

Carbon dioxide and methane are greenhouse gases; sulfur dioxide and nitrogen oxides cause acid rain.

Carbon monoxide is toxic because it binds to haemoglobin and reduces oxygen transport.

Water is tested for purity by boiling point (100 C at 1 atm) and by absence of dissolved solids; distilled water is pure.

Water treatment: filtration removes solids, chlorination kills microorganisms, and boiling is used in the laboratory.

Contact process: sulfur burned to sulfur dioxide, oxidised to sulfur trioxide with a vanadium(V) oxide catalyst, then absorbed in concentrated sulfuric acid.

Fertilisers supply nitrogen, phosphorus and potassium; excess nitrate runoff causes eutrophication of rivers.

Use a burette for accurate variable volumes in titration, a pipette for a fixed volume, and a measuring cylinder for approximate volumes.

Read a meniscus at eye level, taking the bottom of the curve for most aqueous solutions.

Separate mixtures by filtration (insoluble solid), evaporation or crystallisation (soluble solid), and distillation (liquids with different boiling points).

Chromatography separates dissolved substances by differing solubility and adsorption; the Rf value is distance moved by spot divided by distance moved by solvent.

Tests: limewater turns milky with carbon dioxide; a glowing splint relights in oxygen; a lighted splint pops with hydrogen; damp red litmus turns blue with ammonia.

Flame tests: lithium crimson, sodium yellow, potassium lilac, calcium brick red, copper blue-green.

Record measurements to the precision of the instrument, calculate means, and identify anomalous results before plotting graphs.