An atom has a tiny, dense nucleus (containing protons and neutrons) surrounded by electrons in shells. Almost all the mass is in the nucleus, but almost all the volume is empty space where the electrons orbit.
Atomic number (bottom number, or the smaller one) = number of protons, which also fixes the element. Mass number (top number, the bigger one) = protons + neutrons.
Don't mix up mass number and atomic number, and don't forget electrons barely add to the mass but do add to the atom's size.
Electrons fill shells starting from the one closest to the nucleus (lowest energy first). The maximum numbers are 2, then 8, then 8 (for the first 20 elements on this spec).
Elements are arranged in order of increasing atomic number. Columns are groups (share the same number of outer electrons, so similar chemical properties); rows are periods.
Group number tells you outer electrons, not total electrons, and it is the modern table's atomic number order, not atomic weight order, that makes it work.
There are three types of strong chemical bonding: ionic, covalent and metallic. All involve electrons in the outer shell (the highest occupied energy level) rearranging so atoms get a full outer shell, usually matching the nearest noble gas.
Ionic compounds form giant ionic lattices with strong electrostatic forces in all directions. This gives high melting and boiling points. They conduct electricity only when molten or dissolved in water, because the ions are then free to move and carry charge. As a solid the ions are fixed in place, so no conduction happens.
Substances like water, oxygen and CO2 have strong covalent bonds inside the molecule but only weak intermolecular forces between molecules. It is these weak forces that break on melting or boiling, not the covalent bonds, so melting and boiling points are low. Common mistake: students often say the covalent bonds break, they do not. These substances do not conduct electricity because there are no free ions or electrons.
Substances such as diamond, graphite and silicon dioxide have many strong covalent bonds in a giant lattice, giving very high melting points. Diamond: each carbon bonds to 4 others, very hard, does not conduct. Graphite: each carbon bonds to 3 others in layers, with one delocalised electron per carbon, so it conducts electricity and the layers slide over each other (soft, good lubricant).
Metals have giant structures with layers of atoms that can slide over each other, making them malleable. Delocalised electrons allow metals to conduct electricity and heat, both as solids and liquids. Alloys are mixtures of metals (or metal plus another element) with atoms of different sizes, which distort the layers and make alloys harder than pure metals.
Fullerenes (including nanotubes) are carbon molecules that form hollow shapes, used in drug delivery, lubricants and reinforcing materials. State symbols show state at room temperature: (s) solid, (l) liquid, (g) gas, (aq) aqueous (dissolved in water). Particle theory links to state changes and limitations of the simple model (e.g. particles are treated as solid spheres with no forces, which is not fully accurate).
In a chemical reaction, atoms are never created or destroyed, only rearranged. This means the total mass of reactants equals the total mass of products in a closed system. If a reaction seems to lose mass, a gas has probably escaped (e.g. thermal decomposition or a reaction with acid). If it seems to gain mass, a gas from the air (usually oxygen) has been taken in, as in metal oxidation.
Relative atomic mass (Ar) is a weighted average based on isotope abundance. Relative formula mass (Mr) is the sum of the Ar values of all atoms in the formula. Always check the formula carefully — brackets multiply everything inside, e.g. Ca(OH)2 has 1 Ca, 2 O, 2 H.
A mole is 6.02 x 10^23 particles (Avogadro's constant). Moles link mass, particles and gas volume together.
Concentration is usually measured in g/dm3 or mol/dm3.
Remember to convert cm3 to dm3 by dividing by 1000 — this is the single most common exam slip.
Balanced symbol equations show the mole ratio between reactants and products. Use this ratio to calculate reacting masses: find moles of the known substance, use the ratio to find moles of the unknown, then convert to mass using Mr.
The reactant that gets used up first is the limiting reactant — it controls how much product can form, even if the other reactant is in excess. Work out moles of each reactant, compare to the equation ratio, and whichever gives the smaller product amount is limiting.
Titrations find the exact volume of one solution that reacts completely with another, using an indicator to spot the end point. Always do a rough titration first, then repeat until you get concordant results (within 0.10 cm3 of each other), and average only the concordant runs.
Metals can be ranked in the reactivity series: potassium, sodium, lithium, calcium, magnesium, aluminium, (carbon), zinc, iron, (hydrogen), copper, silver, gold.
A more reactive metal displaces a less reactive one from a compound in solution or in an oxide. Carbon and hydrogen sit in the series as non-metal benchmarks: a metal below carbon can be extracted by heating its oxide with carbon (reduction); a metal above carbon needs electrolysis instead.
Oxidation is loss of electrons or gain of oxygen; reduction is gain of electrons or loss of oxygen (OIL RIG). In extraction, the metal oxide is reduced and the carbon is oxidised to carbon dioxide.
Unreactive metals like gold are found native (uncombined) so need no chemical extraction.
Acid + metal makes a salt plus hydrogen gas (test with a lit splint: a squeaky pop confirms hydrogen).
Acid + metal oxide or metal hydroxide makes a salt plus water (neutralisation).
Acid + metal carbonate makes a salt plus water plus carbon dioxide (test with limewater: it turns cloudy/milky).
Hydrochloric acid gives chloride salts, sulfuric acid gives sulfate salts, nitric acid gives nitrate salts.
The pH scale runs 0 to 14. Below 7 is acidic, 7 is neutral, above 7 is alkaline. Universal indicator shows this as a colour gradient; a pH probe gives a precise reading.
Acids release H+ ions in solution; alkalis release OH- ions. Neutralisation is H+ + OH- -> H2O.
To make a soluble salt from an insoluble base: add excess base to the acid, filter off unreacted base, then evaporate/crystallise the salt solution. This ensures all the acid has reacted.
Electrolysis breaks down an ionic compound using electricity, needed when a substance is molten or dissolved (ions must be free to move).
At the cathode (negative electrode), positive ions gain electrons (reduction) - metal is deposited, or hydrogen if the metal is more reactive than hydrogen.
At the anode (positive electrode), negative ions lose electrons (oxidation) - usually a gas like oxygen or chlorine forms.
In electrolysis of concentrated sodium chloride solution: hydrogen at the cathode, chlorine at the anode, leaving sodium hydroxide solution behind.
Exothermic reactions transfer energy to the surroundings, so temperature rises (e.g. combustion, neutralisation). Endothermic reactions take energy in, so temperature falls (e.g. thermal decomposition, some dissolving).
A reaction profile shows the energy of reactants and products; the activation energy is the minimum energy needed to start the reaction.
Bond breaking requires energy (endothermic); bond making releases energy (exothermic). If more energy is released making bonds than used breaking them, the reaction is exothermic overall.
Don't confuse oxidation and reduction directions - always check whether electrons/oxygen are gained or lost.
Remember to filter before evaporating when making a soluble salt from an insoluble base - forgetting this leaves impurities.
At electrodes, always link cathode to reduction and anode to oxidation (an easy memory aid is 'anode = oxidation', both start with vowels).
Rate of reaction measures how fast reactants turn into products. You can measure it two ways: the amount of reactant used up over time, or the amount of product formed over time. Units are usually g/s or cm3/s.
Rate = quantity of reactant used (or product formed) / time taken.
On a graph of product formed against time, the rate at any point is the gradient (steepness) of the tangent to the curve. Steeper = faster. The line flattens out when the reaction finishes because a reactant has run out.
Collision theory explains all of this: reactions only happen when particles collide with enough energy (the activation energy) and the correct orientation.
A reversible reaction can go forwards and backwards, shown with the symbol pointing both ways. In a closed system, a reversible reaction reaches dynamic equilibrium: forward and backward reactions still happen, but at equal rates, so concentrations of reactants and products stay constant.
Le Chatelier's principle: if you change the conditions on a system at equilibrium, the equilibrium shifts to counteract that change.
Common mistake: students think a catalyst shifts the equilibrium position. It does not; a catalyst only speeds up reaching equilibrium, both directions equally.
Crude oil is a mixture of hydrocarbons, mostly alkanes, formed from ancient biomass. Fractional distillation separates it by boiling point: the oil is heated, vapours rise up a fractionating column, and different fractions condense at different heights (short chains at the top/coolest, long chains at the bottom/hottest).
Alkanes are saturated hydrocarbons with the general formula CnH2n+2 (e.g. methane CH4, ethane C2H6, propane C3H8, butane C4H10).
Cracking breaks long-chain, less useful hydrocarbons into shorter, more useful alkanes and alkenes, using heat and a catalyst (catalytic cracking) or steam at high temperature (steam cracking). Alkenes are unsaturated, general formula CnH2n, and contain a C=C double bond. Test for alkenes: bromine water turns from orange to colourless.
Common mistake: mixing up saturated (alkanes, only single bonds) with unsaturated (alkenes, C=C double bond, decolourise bromine water).
Mixtures of metal ions can mask each other's flame colours, so flame tests alone cannot always identify every ion in a mixture.
Roughly 80% nitrogen, 20% oxygen, small amounts of carbon dioxide (about 0.04%), argon and water vapour. This has been roughly the same for the last 200 million years.
1. Early Earth: volcanoes released mostly carbon dioxide and water vapour, plus nitrogen, little or no oxygen.
2. Oceans formed as water vapour condensed; CO2 dissolved into oceans and was locked into carbonate rocks and fossil fuels.
3. Algae and plants evolved and produced oxygen by photosynthesis, gradually building up the oxygen level.
4. Nitrogen built up because it is unreactive, so it stayed in the atmosphere.
Carbon dioxide, methane and water vapour trap outgoing infrared radiation, keeping the planet warm. Human activity (burning fossil fuels, deforestation, agriculture, landfill waste) has increased CO2 and methane levels, correlating with rising global temperatures. This causes rising sea levels, more extreme weather and changes to habitats and species distribution.