An atom has a tiny, dense nucleus (protons + neutrons) surrounded by electrons in shells. The nucleus takes up almost no volume but nearly all the mass. Atoms are roughly 0.1 nanometres (1 x 10^-10 m) across; the nucleus is about 10,000 times smaller than the atom.
Isotopes are atoms of the same element with the same number of protons but a different number of neutrons, so they have different mass numbers. Because they have the same electron arrangement, isotopes react chemically the same way, but their physical properties (like density) can differ slightly. Relative atomic mass (Ar) is a weighted average of an element's isotopes based on their abundance, which is why some Ar values are not whole numbers, like chlorine at 35.5.
Electrons fill shells starting from the one nearest the nucleus, lowest energy first. The maximum numbers are 2, then 8, then 8 (for the first 20 elements). Write configurations like 2,8,1 for sodium. The outer shell electrons control chemical reactivity, which is why elements in the same group behave similarly.
Elements are arranged in order of increasing atomic number. Groups are the vertical columns and share the same number of outer electrons, so they have similar chemical properties. Periods are the horizontal rows, and the period number equals the number of electron shells. Metals sit on the left and centre; non-metals sit on the right, separated roughly by a zig-zag line.
Know the order: Dalton's solid spheres, then Thomson's plum pudding model after discovering the electron, then Rutherford's alpha-particle scattering experiment revealed a small positive nucleus (leading to the nuclear model), then Bohr proposed fixed electron shells, then later work established the neutron (Chadwick).
There are three types of strong chemical bonding: ionic, covalent and metallic.
Ionic compounds form giant ionic lattices. Common mistake: students say ionic bonds are 'weak' because compounds dissolve easily. In fact ionic bonds are very strong, which is why melting and boiling points are high.
There are two types: simple molecular and giant covalent (macromolecular).
Relative atomic mass (Ar) and relative formula mass (Mr) are ratios comparing atoms/compounds to carbon-12, so they have no units.
A mole is a fixed amount of substance containing 6.02 x 10^23 particles (atoms, ions or molecules) - this is Avogadro's constant.
Mass is conserved in a closed system - total mass of reactants equals total mass of products, because atoms are only rearranged, never created or destroyed.
Use balanced symbol equations to find the ratio of moles reacting.
The limiting reactant is used up first and controls how much product forms; the other reactant is in excess.
Concentration is usually measured in g/dm3 or mol/dm3.
Percentage yield = (actual yield / theoretical yield) x 100. It's always less than 100% due to practical losses, incomplete reactions or side reactions.
An exothermic reaction transfers energy to the surroundings, so the temperature goes up. Examples: combustion, neutralisation, most oxidation reactions, and respiration.
An endothermic reaction takes in energy from the surroundings, so the temperature goes down. Examples: thermal decomposition, and citric acid reacting with sodium hydrogencarbonate.
Reversible reactions can be exothermic in one direction and endothermic in the other, with the same amount of energy transferred each way.
In practicals you add reactants to an insulated cup (often inside a beaker of cotton wool for extra insulation) and record the maximum or minimum temperature reached. A lid reduces energy loss to the air, and insulation reduces energy loss through the container. Common mistake: forgetting the lid or insulation, which makes the temperature change look smaller than it really is.
Breaking bonds requires energy (endothermic, an input).
Making bonds releases energy (exothermic, an output).
Overall energy change = energy to break bonds minus energy released making bonds.
If more energy is released making bonds than is used breaking them, the reaction is exothermic overall. If more energy is needed to break bonds than is released, it's endothermic overall.
You may be given bond energy values in kJ/mol and asked to calculate the overall energy change - always show the breaking total, the making total, and the subtraction.
The reactivity series (most to least reactive): potassium, sodium, lithium, calcium, magnesium, aluminium, carbon, zinc, iron, hydrogen, copper, silver, gold.
A more reactive metal displaces a less reactive metal from its compound in solution or oxide - this is a redox reaction.
Oxidation is loss of electrons; reduction is gain of electrons (remember OIL RIG).
In displacement reactions, the more reactive metal is oxidised (loses electrons, forms ions) and the less reactive metal ion is reduced (gains electrons, forms the metal).
Electrolysis breaks down an ionic compound using electricity, needed when a compound can't be reduced by carbon (e.g. aluminium, which is more reactive than carbon).
Positive ions (cations) move to the negative electrode (cathode) and are reduced (gain electrons).
Negative ions (anions) move to the positive electrode (anode) and are oxidised (lose electrons).
In molten ionic compounds, the metal forms at the cathode and the non-metal forms at the anode.
In aqueous solutions, if the metal is more reactive than hydrogen, hydrogen gas forms at the cathode instead of the metal; at the anode, oxygen forms unless a halide ion is present, in which case the halogen forms instead.
Mixing up which electrode is which - cathode is negative, anode is positive.
Forgetting that exothermic means energy OUT to the surroundings (temperature rises), not energy 'created'.
Writing bond energy calculations the wrong way round - it's always bonds broken minus bonds made.
Rate = amount of reactant used up (or product formed) divided by time. Units are usually g/s, cm3/s or mol/s. You can measure rate by tracking mass loss (for gas-producing reactions, on a balance), gas volume (using a gas syringe) or a colour/turbidity change (like the classic sodium thiosulfate + hydrochloric acid 'disappearing cross' experiment).
For a reaction to happen, particles must collide with enough energy - this minimum energy is the activation energy. Anything that increases collision frequency or collision energy speeds up the reaction.
Don't say a catalyst 'speeds up particles' - it lowers the activation energy needed, it does not change particle speed or temperature.
A reversible reaction can go forwards and backwards, shown with the reversible arrow symbol. At dynamic equilibrium (in a closed system) the forward and backward reactions happen at the same rate, so concentrations of reactants and products stay constant - but the reaction has NOT stopped.
Le Chatelier's principle: if you change conditions, the equilibrium shifts to oppose the change.
Crude oil is a mixture of hydrocarbons, mainly alkanes, separated by fractional distillation using differences in boiling point. Shorter chains have lower boiling points, are more flammable and less viscous.
Alkanes: general formula CnH2n+2, saturated (single C-C bonds only). First four: methane (CH4), ethane (C2H6), propane (C3H8), butane (C4H10).
Alkenes: general formula CnH2n, unsaturated (contain a C=C double bond). First three: ethene (C2H4), propene (C3H6), butene (C4H8). Alkenes decolourise orange bromine water - alkanes do not. This is the key test to distinguish them.
Cracking breaks long-chain alkanes into shorter, more useful alkanes and alkenes, using either catalytic cracking (catalyst, around 550C) or steam cracking (high temperature, no catalyst). Cracking meets demand for smaller molecules like petrol and makes alkenes for polymers.
Dip a clean wire in the sample and hold it in a blue Bunsen flame. The colour tells you the metal ion:
Add a few drops of NaOH solution to the unknown solution. A coloured precipitate forms:
Flame emission spectroscopy is a modern instrumental method used instead of flame tests. It is faster, more accurate, more sensitive, and can identify mixtures of ions from a single sample by analysing the exact wavelengths of light emitted, shown as line spectra on a computer.
Earth's early atmosphere was mostly carbon dioxide from volcanic activity, with little or no oxygen. Oceans formed as water vapour condensed, and CO2 dissolved into them, later locking into carbonate rocks and fossil fuels. Algae and plants evolved and produced oxygen through photosynthesis, gradually building up today's atmosphere: roughly 80% nitrogen, 20% oxygen, with small amounts of carbon dioxide, water vapour, and noble gases (mainly argon).
Carbon dioxide, methane, and water vapour are the main greenhouse gases. Human activities (burning fossil fuels, deforestation, farming, and landfill) increase levels of CO2 and methane, enhancing the greenhouse effect and driving global climate change: rising sea levels, melting ice, more extreme weather.