Roadmap

The mole is the SI unit for counting particles, defined so that one mole of any substance contains the same number of entities — which lets a mass measured on a balance be converted into a count of atoms or molecules.

Atoms bond by transferring or sharing electrons, and where a bond sits on the ionic-to-covalent scale is set by the difference in electronegativity — which in turn predicts a substance's melting point, solubility and conductivity.

A balanced equation gives the fixed ratio in which substances react, in moles rather than in grams — so predicting how much product forms means converting to moles, applying the ratio, and converting back.

An acid donates protons and a base accepts them; pH is the negative logarithm of the hydrogen-ion concentration, so each unit of pH is a tenfold change in acidity.

In a redox reaction electrons transfer from one species to another: the donor is oxidised and the acceptor reduced, and because electrons cannot be created the two always happen together in matched amounts.

Ordering the elements by atomic number arranges them into groups whose members share an outer-electron configuration, so properties such as atomic radius, ionisation energy and reactivity vary in predictable trends across a period and down a group.

A reversible reaction reaches equilibrium when forward and reverse rates become equal, so concentrations stop changing while both reactions continue — and a disturbance shifts the balance in the direction that partly offsets it.

Reaction rate depends on how often reactant particles collide with enough energy and the right orientation, so concentration, temperature, surface area and catalysts all change it — while the equilibrium position stays untouched.

A strong electrolyte dissociates essentially completely into ions in solution while a weak electrolyte reaches only partial dissociation equilibrium, and this distinction shows up directly in molar conductivity, which for a strong electrolyte declines only mildly with concentration due to ion-ion interactions but for a weak electrolyte drops sharply as concentration rises because the fraction actually dissociated falls.

Along the boundary between two phases of a pure substance on a phase diagram, the vapour pressure and temperature are locked together by the Clausius-Clapeyron equation, which follows from requiring both phases to have equal Gibbs energy at equilibrium and explains why boiling point rises with pressure and why the boundary's steepness is set by the transition's enthalpy.

Raoult's law states that an ideal solution's vapour pressure from each volatile component is that component's pure vapour pressure scaled by its mole fraction, and this simple dilution effect is the common origin of the colligative properties — boiling point elevation, freezing point depression, and osmotic pressure — which depend only on how many solute particles are dissolved, not on what they chemically are.

The Nernst equation relates an electrochemical cell's actual voltage to its standard potential and the concentrations (activities) of the species involved, showing that cell potential is not a fixed constant but shifts with concentration in a way that reaches zero exactly when the cell reaction reaches equilibrium.

A reaction's rate law expresses its rate as a product of a rate constant and the concentrations of reactants raised to experimentally determined powers — the reaction order in each species — which must be measured, not read off the balanced equation, since a reaction's order reflects its actual step-by-step mechanism rather than its overall stoichiometry.

The Arrhenius equation, k = A·e^(−Ea/RT), captures why reaction rates rise sharply with temperature: only molecules colliding with kinetic energy above the activation energy Ea can react, and the exponentially growing fraction of molecules meeting that threshold as temperature rises is what drives the rate constant's strong temperature dependence.

A catalyst speeds up a reaction by opening an alternative reaction pathway with a lower activation energy, without being consumed itself and without changing the reaction's overall equilibrium position, and is classified as homogeneous when it shares the same phase as the reactants (as with enzymes in solution) or heterogeneous when it exists in a separate phase, most often as a solid surface that reactant molecules adsorb onto.

A mass spectrometer identifies molecules by ionising a sample, separating the resulting ions by their mass-to-charge ratio, and recording the abundance of each ratio, producing a mass spectrum whose fragmentation pattern acts as a structural fingerprint of the original molecule in addition to giving its precise molecular mass.

A coloured or UV-absorbing solution absorbs light most strongly at wavelengths matching electronic transitions in its molecules, and the Beer-Lambert law states that absorbance is directly proportional to both the absorbing species' concentration and the path length of light through the sample, turning a simple absorbance measurement into a direct quantitative determination of concentration.

A functional group is the specific arrangement of atoms in a molecule that reacts, so molecules with the same functional group undergo the same characteristic reactions regardless of how large or different the rest of the molecule is.

A covalent bond can break homolytically, with each atom keeping one electron to form two neutral radicals, or heterolytically, with both electrons going to one atom to form a cation and an anion, and which pathway a reaction follows determines whether the reactive species produced are radicals or the electrophiles and nucleophiles that dominate ionic organic chemistry.

An aromatic ring like benzene has its electrons delocalised evenly around the whole ring rather than fixed in alternating single and double bonds, which makes it far more stable than a normal alkene and changes how it reacts.

The relatively weak, exposed pi bond in an alkene or alkyne is broken open by an addition reaction, in which an electrophile is attracted to the electron-rich pi bond, forms a bond to one carbon, and generates a cationic intermediate that a nucleophile then attacks, converting the double or triple bond into new single bonds to the added atoms.

The carbon-oxygen double bond in a ketone or aldehyde is strongly polarised, leaving the carbon electron-poor and directly attractive to nucleophiles, which add across the pi bond to form a tetrahedral alkoxide intermediate — the same polarisation also lets the alpha carbon's C-H bonds ionise as a weak acid, opening a second, complementary reactivity mode called enolisation.

Benzene's delocalised ring of electrons reacts with electrophiles by substitution rather than addition, because the intermediate formed by attack keeps a partial ring and the final step re-forms the aromatic system by losing a proton rather than gaining a second group, preserving the large stabilisation energy that addition would otherwise destroy.

Carboxylic acid derivatives (acid chlorides, anhydrides, esters, and amides) share a carbonyl carbon bonded to a leaving group, and a nucleophile attacking that carbon forms the same tetrahedral intermediate seen in simple carbonyl addition, except the leaving group is then expelled to restore the carbonyl — an addition-elimination sequence whose overall rate is set by how good a leaving group each derivative carries.

A nucleophile is an electron-rich species that attacks an electron-poor carbon, displacing a leaving group; the reaction can happen in one concerted step or in two steps through a stable intermediate, and which pathway operates changes what predicts the outcome.

Molecular orbital theory builds bonding by combining atomic orbitals into new orbitals delocalised over the whole molecule, with bonding combinations lower in energy than the original atomic orbitals and antibonding combinations higher, while orbital hybridisation is a complementary bookkeeping device that mixes an atom's own orbitals into new shapes matching the geometry actually observed, such as the tetrahedral sp3 arrangement around a typical carbon atom.

An element's oxide is basic if the element is strongly metallic, acidic if the element is strongly nonmetallic, and amphoteric for elements near the metal-nonmetal boundary, and the same trend continues into oxoacids — acids containing a central atom bonded to oxygen — whose acid strength increases with the electronegativity and oxidation state of that central atom.

Hydrogen forms compounds across almost the entire periodic table, ranging from ionic hydrides with reactive metals to covalent hydrides with nonmetals, and passing electric current through water or a molten salt (electrolysis) drives the otherwise non-spontaneous decomposition reaction by forcing electrons onto cations at the cathode and pulling them from anions at the anode.

The electrochemical (activity) series ranks metals by their tendency to lose electrons and oxidise, predicting both which metal in a pair will corrode preferentially when they are in electrical contact and why a more reactive 'sacrificial' metal can protect a less reactive one from corrosion by oxidising in its place.

Group 14 spans carbon, whose ability to form strong bonds with itself and a wide range of other elements underlies the whole of organic chemistry, through silicon, whose comparably strong bonds to oxygen instead make silicate minerals the dominant chemistry of the solid Earth, down to increasingly metallic germanium, tin, and lead.

The halogens are the most reactive non-metal group because each atom needs only one more electron for a full outer shell, and their oxidising power falls steadily going down the group as that electron gets harder to attract.

Group 15 spans from nitrogen, a diatomic gas whose triple bond makes it exceptionally unreactive, through phosphorus and arsenic to the increasingly metallic antimony and bismuth, and nitrogen's own chemistry is dominated by an unusually wide range of oxidation states, from -3 in ammonia to +5 in nitric acid, reflecting how many different bonding arrangements its five valence electrons can support.

Transition metals fill an inner d-subshell rather than the outermost one, which lets them adopt several oxidation states, form coloured compounds, and bind surrounding molecules as coordination complexes held together by donated electron pairs.

The Langmuir isotherm models adsorption of a gas onto a solid surface as a dynamic equilibrium between molecules sticking to and leaving a finite number of identical surface sites, predicting that surface coverage rises linearly with pressure at low pressure but saturates as every site fills at high pressure, unlike the unbounded, ever-increasing coverage a simple proportional model would wrongly predict.

A colloidal particle suspended in a liquid picks up a surface charge, which attracts a cloud of counter-ions into a diffuse electric double layer around it; when the particle moves relative to the liquid — under an applied field (electrophoresis) or under gravity — only part of that cloud moves with it, so the electrokinetic (zeta) potential at the slipping plane, not the bare surface charge, is what actually governs how strongly like-charged particles repel each other and therefore whether the dispersion stays stable or coagulates.