PMT CHEMISTRY DEFINITIONS

Atomic Number

the number of protons in the nucleus of an atom.

Electron

a negatively charged subatomic particle which orbits the nucleus at various energy levels. The relative mass of an electron is 1/1836.

Ion

a charged atom or molecule.

Isotopes

atoms of the same element with the same number of protons and electrons but different numbers of neutrons. Isotopes of an element have different masses.

Mass Number

the total number of protons and neutrons in the nucleus of an atom.

Mass Spectrometry

an instrument which gives accurate information about relative isotopic mass and the relative abundance of isotopes.

Neutron

a neutral subatomic particle found in the nucleus of an atom. The relative mass of a neutron is 1.

Proton

a positively charged subatomic particle found in the nucleus of an atom. The relative mass of a proton is 1.

Relative Abundance

the amount of one substance compared with another.

Relative Atomic Mass

the weighted mean mass of an atom compared with 1/12th mass of an atom of carbon-12.

Relative Isotopic Mass

the mass of an atom of an isotope compared with 1/12th mass of an atom of carbon-12.

Relative Formula Mass

the mass of the formula unit of a compound with a giant structure. For example, NaCl has a relative formula mass of 58.44 g mol-1 .

Ammonium ion

an ion with the formula NH4+

Hydroxide

an ion with the formula OH-

Ionic Compound

a compound which is made up of oppositely charged ions that are held together by electrostatic forces. These compounds generally have high melting and boiling points. Typically, ionic compounds are soluble and can conduct electricity when liquid or aqueous (but not when solid).

Nitrate

an ion with the formula NO3-

Silver ion

has the formula Ag+

State symbols

symbols within a chemical equation which indicate the state of each compound under the reaction conditions. (g) gaseous, (l) liquid, (s) solid and (aq) aqueous.

Sulfate

an ion with the formula SO42-

Zinc ion

has the formula Zn2+

Amount of substance

the quantity that has moles as its units, used as a way of counting atoms. The amount of substance can be calculated using mass (n = m/M), gas volumes (n = pV/(RT)) or solution volume and concentration (n = CV).

Anhydrous

a crystalline compound containing no water.

Atom Economy

a measure of the amount of starting materials that end up as useful products. A high atom economy means a process is more sustainable as there is less waste produced.

Percentage atom economy

(molecular mass of desired product/ sum of molecular masses of all reactants) x 100

Avogadro Constant (NA)

the number of particles per mole of substance (6.02 x 1023 mol-1).

Composition by mass

the relative mass of each element in a compound.

Hydrated

a crystalline compound that contains water.

Ideal Gas

a gas which has molecules that occupy negligible space with no interactions between them.

The ideal gas equation is

pV = nRT.

Molar Gas Volume

the volume of 1 mole of gas (units: dm3 mol-1).

Molar Mass

mass per mole of a substance (units: g mol-1).

Mole (mol)

the amount of any substance containing as many particles as there are carbon atoms in exactly 12g of carbon-12 isotope.

Percentage yield

(actual yield/ theoretical yield) x 100

Relative Molecular Mass (Mr)

the average mass of one molecule of an element or compound compared to 1/12th the mass of an atom of carbon-12.

Stoichiometry

the relative quantities of substances in a reaction.

Water of Crystallisation

water molecules that form part of the crystalline structure of a compound.

Acid

compounds that release H+ ions in aqueous solution. Common acids include: HCl, H2SO4, HNO3 and CH3COOH.

Alkali

water soluble bases. Alkalis release OH- ions into aqueous solution. Common alkalis include NaOH, KOH and NH3

Neutralisation

a reaction between H+ and OH-, forming water. This may be a reaction between an acid and a base to form a salt (types of bases include carbonates, metal oxides and alkalis).

Strong Acid

an acid that completely dissociates in solution.

Weak Acid

an acid that only partially dissociates in solution

Oxidation Number

a number that represents the number of electrons lost or gained by an atom of an element. A positive oxidation number indicates the loss of electrons. Roman numerals are typically used to indicate the oxidation number of elements that may have different oxidation states (e.g. iron(II) and iron(III)).

Atomic Orbital

a region of space around the nucleus that can hold up to 2 electrons with opposite spins. There is 1 orbital in the s subshell, 3 orbitals in the p subshell and 5 orbitals in the d subshell. Orbitals are filled in order of increasing energy, with orbitals of the same energy occupied singly before pairing.

Electronic Configuration

the arrangement of electrons into orbitals and energy levels around the nucleus of an atom / ion.

Energy Level

the shell that an electron is in.

Shell

the orbit that an orbital is in around the nucleus of an atom. The shell closest to the nucleus is the first shell. The outermost shell that is occupied by electrons is the valence shell.

Sub-shell

a subdivision of the electronic shells into different orbitals. The types of subshell are s, p, d and f.

Average bond enthalpy

the average energy required to break a bond, used as a measurement of the strength of a covalent bond. The average bond enthalpy is measured using a variety of molecules that contain a specific bond.

Bonding pair

a pair of outer-shell electrons involved in bonding.

Covalent bond

a strong bond formed between 2 atoms due to the electrostatic attraction between a shared pair of electrons and the atomic nuclei.

Dative Covalent (Coordinate) bond

a type of covalent bond in which both of the electrons in the shared pair come from one atom.

Electronegativity

the ability of an atom to attract bonding electrons in a covalent bond. This is often quantified using Pauling's electronegativity values. Electronegativity increases towards F in the periodic table.

Electron Pair Repulsion Theory

pairs of electrons around a nucleus repel each other so the shape that a molecule adopts has these pairs of electrons positioned as far apart as possible. Lone pairs offer more repulsion than bonding pairs as they are closer to the nucleus of the central atom.

Hydrogen Bonding

a type of intermolecular bonding that occurs between molecules containing N, O or F and a H atom of -NH, -OH or HF. A lone pair on the electronegative atom (N, O or F) allows the formation of a hydrogen bond.

Intermolecular Forces

interactions between different molecules. Types of intermolecular forces including permanent dipole-dipole interactions and induced dipole-dipole interactions (both of these are also known as van der Waals' forces) as well as hydrogen bonding.

Ionic Lattice

a giant structure in which oppositely charged ions are strongly attracted in all directions.

Linear

the shape of a molecule in which the central atom has 2 bonding pairs.

London (Dispersion) Forces

induced dipole-dipole interactions caused when the random movement of electrons creates a temporary dipole in one molecule which then induces a dipole in a neighbouring molecule.

Lone Pair

a pair of outer-shell electrons not involved in bonding.

Macroscopic Properties

properties of a bulk material rather than the individual atoms/ molecules that make up the material.

Non-linear

the shape of a molecule in which the central atom has 2 bonding pairs and 2 lone pairs.

Permanent Dipole

a permanent uneven distribution of charge.

Polar Molecule

a molecule that contains polar bonds with dipoles that don't cancel out due to their direction (must be unsymmetrical).

Pyramidal

the shape of a molecule in which the central atom has 3 bonding pairs and 1 lone pair.

Simple Molecular Lattice

a solid structure made up of covalently bonded molecules attracted by intermolecular force (e.g. I2 and ice). These compounds generally have relatively low melting and boiling points and are typically insoluble in water but soluble in organic solvents. Molecular substances don't conduct electricity.

Trigonal bipyramidal

the shape of a molecule in which the central atom has 5 bonding pairs. Trigonal Planar: the shape of a molecule in which the central atom has 3 bonding pairs.

Atomic (Proton) Number

​ the number of protons in the nucleus of an atom.

Bohr Model

​describes an atom as a small dense nucleus with electrons orbiting around the nucleus. This model explains different periodic properties of atoms.

d-block

​the part of the periodic table in which the elements have their highest energy electron in a d-orbital.

Giant Covalent Lattice

​a network of atoms bonded by strong covalent bonds (e.g. carbon (diamond, graphite and graphene) and silicon). Giant covalent lattices typically insoluble with a high melting and boiling point due to the presence of strong covalent bonds. They are also poor electrical conductors as they don't contain mobile charged particles.

Giant Metallic Lattice Structure

the structure of all metals, made up of cations and delocalised electrons. Giant metallic structures are typically insoluble with a high melting and boiling points due to strong electrostatic forces of attraction between cations and electrons. Metals are good electrical conductors due to the presence of delocalised electrons (mobile charges).

Group

​ a column in the periodic table.

Melting Point

t​ he temperature at which a solid melts and becomes a liquid. This increases from giant metallic to giant covalent structures then decreases to simple molecular structures.

Metallic Bonding

​strong electrostatic attraction between cations and delocalised electrons.

p-block

​the part of the periodic table in which the elements have their highest energy electron in a p-orbital.

Period

​a row in the periodic table.

Periodicity

​ a repeating trend in physical and chemical properties across the periods of the periodic table.

s-block

​the part of the periodic table in which the elements have their highest energy electron in an s-orbital.

Successive Ionisation Energies

​the energy required to remove each electron one-by-one from one mole of gaseous atoms / ions.

Base

​a substance that can accept H+​ ​ions from another substance. Group 2 compounds can be used as bases: Ca(OH)2​ ​is used to neutralise acidic soils in agriculture and Mg(OH)2​ ​and CaCO3​ ​are used as antacids to treat indigestion.

First Ionisation Energy

​the removal of one mole of electrons from one mole of gaseous atoms. Factors which affect the first ionisation energy are: the strength of attraction between the electron and the nucleus, the nuclear charge and the atomic radius.

Group 2 Oxide

​a compound with the general formula MO, where M is a group 2 element. When group 2 oxides react with water, they form an alkaline solution, with alkalinity increasing down the group.

Redox

​ a reaction in which oxidation of one element and reduction of another occurs. During a redox reaction involving group 2 elements, 2 electrons are lost to form 2+ ions. Group 2 elements undergo redox reactions with water, oxygen and dilute acids. During a redox reaction involving group 7 elements, 1 electron is gained to form 1- ions.

Second Ionisation Energy

​the removal of one mole of electrons from one mole of gaseous 1+ ions to form one mole of 2+ ions.

Boiling Point

​the temperature at which a liquid boils and becomes a gas. Boiling point increases down group 7 due to the increasing strength of London Forces between the halogen molecules.

Diatomic Molecules

​molecules that are made up of 2 atoms. Halogens are diatomic.

Displacement Reaction

​a reaction in which one atom is replaced by another. Halogens can undergo displacement reactions as their reactivity decreases down the group. The more reactive halogen will displace the less reactive halogen from a solution of its salt.

Disproportionation

​the oxidation and reduction of the same element. Examples include the water treatment (reacting chlorine with water) and bleach formation (reacting chlorine with cold, dilute aqueous sodium hydroxide).

Electron Configuration

​the arrangement of electrons into orbitals and energy levels around the nucleus of an atom / ion. The halogens have a s2​ p​ 5​ ​outer shell electron configuration. Group 2 elements have an S2 outer shell electron configuration

Induced Dipole-Dipole Interactions

​forces of attraction between molecules caused when the random movement of electrons creates a temporary dipole in one molecule which then induces a dipole in a neighbouring molecule.

Precipitation Reaction

​a reaction in which two aqueous solutions are combined to form an insoluble salt (a precipitate). Halide anions undergo precipitation reactions with aqueous silver ions.

Water Treatment

​ the addition of chlorine to water to kill bacteria. The risks associated with the use of chlorine to treat water are the hazards of toxic chlorine gas and the possible risks from the formation of chlorinated hydrocarbons.

Ammonium Ion

​an ion with the formula NH4​ +​ .​ The test for ammonium ions is a reaction with warm NaOH, which forms NH3​ .​

Cation

​ a positively charged ion.

Carbonate

​a salt containing the CO​ 2-​ anion. A reaction between a carbonate and H+​ ​will form CO2​(​g).

Halide

​a salt containing a group 7 anion. Cl-​ ,​ Br-​ ​and I-​ ​can be tested for using a solution of silver ions as this reaction forms a coloured precipitate. The solubility of the precipitate is then tested using dilute and concentrated ammonia.

Sulphate

​a salt containing the anion SO​ 2-.​ A reaction between SO​ 2- a​ nd Ba2​ +(​ aq) will form a precipitate.

Activation Energy

​the minimum energy required for a reaction to take place.

Average Bond Enthalpy

​the energy required to break one mole of gaseous bonds. Actual bond enthalpies may differ from the average as the average bond enthalpy considers a particular bond in a range of molecules.

Endothermic

​a reaction which takes in energy from surroundings (ΔH is positive). More energy is required to break bonds than is released by making bonds. The energy of the products is higher than the reactants.