Notes

Atom - Positively charged nucleus containing most of the mass, surrounded by atomic shells with orbiting electrons of negative charge and negligible mass.
Atomic number - Defines the element. The atomic number of the nucleus, also the proton number, shows the number of protons in the nucleus (Unless it’s a charged ion, it’s also the number of electrons orbiting an atom)
Mass number - The mass of the atom, the number of protons added to the number of neutrons (Mass of a proton and a neutron are both 1, the mass of an electron, 1/2000, is negligible).
Example - On the periodic table Carbon has an atomic number of 6 and a mass number of 12, this means that Carbon has 6 protons, 6 neutrons (Number of protons + number of neutrons = mass number, therefore the mass number - number of protons = number of neutrons) and 6 electrons.
Isotope - An atom with the same number of protons but a different number of neutrons.
Example - Carbon also exists as Carbon-14, which has a mass number of 14. The number of protons and electrons are unchanged but the number of neutrons is 8.
Relative atomic mass - Average weighted mass of an atom compared with 1/12th of the mass of an atom of Carbon-12.
Relative isotopic mass - Mass of an atom of an isotope compared with 1/12th of the mass of an atom of Carbon-12.
Amount of substance - The number of atoms, has the unit ‘mole’ (mol).
Avogadro’s constant - The number of particles per mole (6.02*10^23 mol^-1)
Molar mass - Mass per mole of a substance, has the unit grams per mole (g mol^-1)
Example - Carbon dioxide, CO2, has a molar mass of 44 g mol^-1 as Carbon has a mass of 12 and Oxygen has a mass of 16, there are two Oxygen atoms -> 12 + (16*2) = 44.
Empirical formula - The simplest whole number ratio of atoms of each element present in a compound
Molecular formula - The actual number of atoms of each element in a molecule
Example - The molecular formula for Glucose is C6H12O6 but the empirical formula is CH2O as the number of all the elements in C6H12O6 are divisible by 6.
Concentration - The amount of solute (in mol) per 1dm^3 of solution, has the units moles per dm^3 (mol dm^-3). A concentrated solution has a high amount of solute per dm^3, a dilute solution has a low amount of solute per dm^3.
Acids - Proton donors. When they’re in water they release H+ ions into the solution.
Examples - HCl (Hydrochloric acid), HNO3 (Nitric acid), H2SO4 (Sulfuric acid), H3PO4 (Phosphoric acid).
Bases - Proton acceptors, they take H+ ions and neutralise acids.
Examples - Metal oxides such as MgO (Magnesium oxide) and CuO (Copper Oxide).
Metal hydroxides such as NaOH (Sodium hydroxide), KOH (Potassium hydroxide) and Mg(OH)2 (Magnesium hydroxide).
Ammonia (NH3) and Amines such as CH3NH2 (Methylamine).
Alkali - Soluble base which releases OH- ions in solution.
Examples - Metal hydroxides such as those listed above and Ammonia.
Salts - Produced when a H+ ion is replaced by a metal ion or NH4 +.
Examples - 2HCl + CuCO3 -> CuCl2 + CO2 + H2O. Where the H+ ion from HCl has been replaced by a metal ion (Cu 2+), forming the salt, CuCl2 (Copper chloride).
H2SO4 + NaOH -> Na2SO4 + H2O. Where the H+ ions in H2SO4 have been replaced by a metal ion (Na-), forming the salt Na2SO4 (Sodium sulfate).
HNO3 + NH3 -> NH4NO3. Where the H+ has datively bonded to NH3 (Ammonia) to form an NH4 + ion which bonds with the NO3 - ion to form the salt NH4NO3.
Anhydrous - Substance containing no water molecules.
Hydrated - Crystalline compound containing water molecules.
Water of crystallisation - Water molecules which form part of the crystalline structure of a (hydrated) compound.
Oxidation - Loss of electrons (OIL, oxidation is loss), increase in oxidation number shows that a species has been oxidised.
Reduction - Gain of electrons (RIG, reduction is gain), decrease in oxidation number shows that a species has been reduced.
Roman numerals (I, II, III, IV, V etc) - The numbers state the oxidation state of an ambiguous element in a molecule.
Example - Sodium chlorate exists in a few different oxidation states, one is NaClO. In this, O has an oxidation number of -2, Na has an oxidation number of +1. To have an overall oxidation number of 0, Cl must have an oxidation number of +1, giving this Sodium chlorate the systematic name Sodium chlorate (I).
Sodium chlorate also exists as NaClO3. O has an oxidation number of -2, multiplied by 3 as there are 3 Oxygen atoms (-6), and Na has an oxidation number of +1. To have an overall oxidation number of 0, Cl must have an oxidation number of +5, giving this Sodium chlorate the systematic name Sodium chlorate (V).
I’m not quite certain why chlorine is the ambiguous element and sodium isn’t, but my guess is that since Sodium has 1 electron in its outer shell and Chlorine has 7, it’s easier to remove more electrons from Chlorine to give it more oxidation states.
Redox reaction - A reaction in which species are both reduced and oxidised.
Disproportionation reaction - A reaction in which the same species is both reduced and oxidised.
First ionisation energy - The energy required to remove one electron from the outer shell from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
Second ionisation energy (And so on) - The energy required to remove one electron from the outer shell from one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions.
Nuclear charge - The attraction from the protons in the nucleus with electrons. Protons have a positive charge and electrons have a negative charge, the greater the number of protons (The atomic number of the element), the greater the nuclear charge. Increases along a period. A higher nuclear charge makes it more difficult to remove an electron from the atom’s outer shell, causing the ionisation energy to increase.
Electron shielding - The number of shells of electrons between the nucleus and the outer shell of electrons. Increases down a group. Electrons repel each other so the more electrons repelling the electrons in the outer shell, the easier it’ll be to remove them, causing the ionisation energy to decrease.
Atomic radius - The distance from the outer shell of electrons to the nucleus of the atom. Increases down a group. A greater atomic radius means the outer shell electrons are under less influence from the attraction from the nucleus, this makes it easier to remove the electrons, causing the ionisation energy to decrease.
Orbital - A region that can hold up to two electrons of opposite spins (Up and down). Orbitals in an s-subshell are spherical, orbitals in a p-subshell are hourglass shaped.
Subshell - The space an electron can occupy within each shell (218:32). An s-subshell has one orbital (And so can hold a total of 2 electrons), a p-subshell has 3 orbitals (6 electrons), a d-subshell has 5 orbitals (10 electrons), a f-subshell has 7 orbitals (14 electrons).
Periodic table ‘blocks’ - The blocks in a periodic table show which subshell the outer shell electron lies in. s-block is groups 1 and 2 (With Helium), p-block is groups 3 to 0 (Without Helium), d-block is the transition metals.
Examples - Na, the first element in group 1, on the third row of the periodic table, is in the s-block. It has an electron configuration of 1s2, 2s2, 2p6, 3s1.
Sc, a transition metal, is in the d-block, it has an electron configuration of 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d1 (The 4s subshell is at a lower energy than the 3d subshell so it fills up first - After the 3d subshell is filled, 4p fills up)
C, the second element in group 4, on the second row of the periodic table, is in the p-block. It has an electron configuration of 1s2, 2s2, 2p2.
Ionic bonding - Electrostatic attraction between oppositely charged ions, where electrons are given and taken.
Covalent bonding - A bond formed when atoms share pairs of electrons.
Dative(/coordinate) covalent bonding - A bond formed when one of the bonding atoms gives both of a pair of electrons.
Bonding pairs (of electrons) - The number of pairs of electrons in a bond.
Lone pairs (of electrons) - The number of pairs of electrons that aren't part of a bond.
Bond angle - The angle between bonds in a molecule, which are defined by the number of bonding pairs of electrons and lone pairs of electrons as the electrons repel each other and space out evenly. Lone pairs repel more than bonding pairs.
Examples - CO2 has two bonding pairs (2 oxygens to carbon) and no lone pairs on the central atom (carbon), giving it a bond angle of 180 degrees (Called 'Linear').
H2O has two bonding pairs of electrons around the central atom (2 hydrogens bonding to an oxygen) and lone pairs on the oxygen, this gives it a bond angle of 104.5 degrees (Called 'Non-linear').
BF3 has 3 bonding pairs and no lone pairs on boron, bond angle of 120 degrees ('Trigonal planar').
NH3 has 3 bonding pairs (3 hydrogens to nitrogen) and lone pairs on the nitrogen, giving it a bond angle of 107 degrees (Called 'Pyramidal').
CH4 has 4 bonding pairs and no lone pairs on carbon, bond angle of 109 degrees ('Tetrahedral').
SF6 has 6 bonding pairs and no lone pairs on sulfur, has a bond angle of 90 degrees ('Octahedral').
Intermolecular forces - The forces of attraction between molecules. The strength or amount of intermolecular forces are what affect the substance's melting/boiling points - Stronger intermolecular forces require more (heat) energy to break.
Electronegativity - The ability of an atom to attract the bonding electrons in a covalent bond
An atom's electronegativity is defined by the number of electrons it has compared to the number of electrons the bonded atom has. If an atom is more electronegative than another, it'll have a greater attraction with the bonding pairs, this gives the atoms a small charge difference (Small charge differences are shown by delta-positives(d+) and delta-negatives(d-)). This is called a permanent dipole - as the electrons are not evenly distributed the bond between the atoms is now polar.
Dipole-Dipole interactions - An intermolecular force. Molecules with permanent dipoles allow for weak intermolecular bonds to be formed between the molecules - The d+ of one atom attracts the d- of an atom in another molecule, and that molecule does the same thing to other molecules and so on.
Example - H-Cl (HCl, Hydrochloric acid) has a d+ hydrogen and a d- chlorine
H-Cl d- ------- d+ H-Cl
Van der Waals forces - Van der Waals forces are another type of intermolecular force. Van der Waals are formed when the movement of electrons unbalances the distribution of the charge in the shells, this causes an instantaneous dipole to form - This in turn attracts/repels electrons in neighbouring molecules which allows them to form instantaneous dipoles (The dipoles formed in other molecules as a result of an instantaneous dipole are called 'induced dipoles') and so on.
Van der Waals forces increase as the number of electrons in the molecules increases - This is why the molecules of the halogens (Which all exhibit Van der Waals forces), F2, Cl2, Br2 etc. become more solid down the group, at room temperature and pressure (RTP), F2 is a gas and I2 is a solid.
Hydrogen bond - A hydrogen bond is a strong dipole-dipole interaction between a hydrogen atom (Which is electron deficient) and a lone pair of electrons on a highly electronegative atom (Usually oxygen or nitrogen). The hydrogen is d+ and the other atom must be a d-.
Example - H-O-H (H2O, water) has an oxygen with two lone pairs of electrons, the d+ hydrogen in other water molecules would attract to the d- of the oxygen.
H2-O d- ------ d+ H-O-H
Hydrogen bonds are the cause of water's many anomalous properties including the fact that ice has a lower density than water (Ice has an open lattice with rigid hydrogen bonds holding the molecules apart, after it has melted the bonds are less rigid and allow the molecules to be closely packed together) and the fact that water has a relatively high melting and boiling point (Strong bonds require more energy to break, more heat is necessary to give that energy).
In terms of strength of intermolecular forces, hydrogen bonds are generally the strongest, followed by dipole-dipole interactions and Van der Waals is the weakest.
Metallic bonding - The attraction of positive ions to delocalised electrons (Metals have 'seas' of delocalised electrons allowing them to form many metallic bonds).
Giant ionic lattice - Formed by the attraction of oppositely charged ions, each ion is surrounded by the oppositely charged ions and the ions attract each other to form a giant lattice.
Example - NaCl (Sodium chloride) is an ionic compound, it can form giant ionic lattices, the Na+ can attract Cl- from other NaCl molecules and the Cl- can attract Na+, this leads to the Na+ being surrounded by Cl- and vice versa.
High melting and boiling points as they are held together by strong electrostatic forces which require a lot of energy to break (the greater the charge difference the more energy required to break the bonds).
Doesn't conduct electricity when solid but does conduct electricity when molten as the ions cannot move in the solid ions and so cannot carry the charge, however the ions are free to move when molten.
Can dissolve in polar solvents such as water as the water molecules surround the ions, forming a solution.
Giant covalent lattice - Three dimensional structure of atoms bonded together by strong intramolecular covalent bonds.
Example - Diamond is made up completely of carbon atoms, each carbon is covalently bonded to four other carbon atom which are bonded to four other carbon atoms and so on. The many strong covalent bonds are why diamond is so tough to break.
High melting and boiling point as high energies are required to break the strong covalent bonds.
Not conductors of electricity as there are no free charged particles (With the exception of Graphite as there are delocalised electrons between the layers)
Insoluble in polar and non-polar solvents as the covalent bonds are too strong to be broken by both solvents.
Giant metallic lattice - Lattices which contain ionised atoms in fixed positions with delocalised outer shell electrons which spread, and can freely move, throughout the structure.
High melting/boiling point as the attraction between ions and electrons is strong and so a lot of energy is required to dislodge the ions.
Good electrical conductivity as the ions can move freely around the lattice and so can carry the charge.
Malleable and ductile as the delocalised electrons can move, giving the structure a degree of 'give' which allows atoms to slide past each other.
Simple molecular lattice - Three dimensional structure of molecules bonded together by weak intermolecular forces, such as Van der Waals forces.
Low melting and boiling points as the weak forces between the molecules require little energy to break them.
Cannot conduct electricity as there are no charged particles to move around.
Soluble in non-polar solvents as Van der Waals forces form between the solvent and the simple molecular structures.
Periodicity - A repeating pattern across different periods (rows).
Example - Atoms of elements in a group have similar outer shell electron configurations and so have similar physical and chemical properties.