IB Chemistry glossary of terms and definitions

(1) a process in which one substance penetrates into the interior of another material, eg. water is absorbed into a sponge, or a gas is absorbed by a liquid. The process may be purely physical, or may involve chemical bonding (see also adsorption);

(2) the process in which energy enters a substance and brings about change, eg. electromagnetic radiation in the visible region may bring about electronic transitions, or infra-red radiation may cause the vibration of bonds.

Acid anhydride

This term is used to describe

(1) a non-metal oxide that reacts with water to give an acidic solution, eg. sulfur trioxide is the acid anhydride of sulfuric acid;

(2) an organic acid derivative containing the anhydride functional group (-CO-O-CO-). An acid anhydride reacts with water to give the acid or acids from which it is derived, eg. ethanoic anhydride, CH₃-CO-O-CO-CH₃ gives ethanoic acid on reaction with water.

Acid ionisation constant, Ka

The equilibrium constant for the ionisation of a weak acid, also known as the acid dissociation constant.

Acid salt
An acid salt is a compound derived from a polyprotic acid where at least one of the acid hydrogen atoms in the acid molecule is replaced by a metal, while some ‘acid’ hydrogen still remains, eg. sodium hydrogensulfate is an acid salt derived from sulfuric acid.

Activated complex – see transition state.

Activity

This term is used to describe

(1) the capacity of a substance to undergo chemical change. It can be used to indicate that a substance undergoes numerous reactions, and it may also imply that the reactions are vigorous;

(2) the activity of a radioactive substance is the number of atoms of that substance that disintegrate per unit time;

(3) the effective concentration of a substance in a solution. Activity values are used in thermodynamic calculations to allow for the non-ideality of real systems.

Activity series – see reactivity series

Allotrope
An allotrope is a form of an element that exists in the same state as other forms of the element at a given temperature and pressure. Different allotropes of the same element have different physical properties, but usually have very similar chemical properties. Examples of elements that form allotropes are oxygen (with allotropes dioxygen and trioxygen in the gas state), and carbon (with allotropes diamond, graphite and fullerine in the solid state).

Amount
This term should be reserved for use when amount of substance is intended. In other cases the term quantity may be used, eg. a quantity of charge may be measured in coulombs but the amount of electrons carrying the charge would be measured in moles.

Amount of substance

The term is a quantity that is proportional to the number of particles in a given sample of a substance. The particles may be atoms, ions, molecules, electrons, etc., but should be specified. The amount of substance may also refer to combined particles of different types, as indicated by a formula, eg. the amount of substance of magnesium chloride, MgCl₂, would indicate a value proportional to both magnesium ions and chloride ions, with the amount of chloride ions being twice that of the magnesium ions. The amount of substance is measured in units of moles.

Ampholyte - see amphoteric (2)

Amphoteric

This term is used

(1) to describe an oxide or hydroxide that is a water-insoluble metal oxide or hydroxide, which may exhibit properties both of an acid and a base. It may behave like an acid in that it reacts with a strong alkali to give a salt, but it may also behave like a base in that it reacts with an acid to give a (different) salt, eg. the amphoteric oxide, aluminium oxide, reacts with sodium hydroxide to give sodium aluminate, and with hydrochloric acid to give aluminium chloride;

(2) to refer to a water soluble substance, such as an amino acid, that can behave both as an acid or a base. Such a substance is an electrolyte and may be referred to as an ampholyte.

Anode

The electrode at which oxidation (ie. loss of electrons) takes place.

Atomic mass unit (symbol u)

A unit of mass equal to exactly one-twelfth of the mass of the carbon-12 atom, sometimes also called the dalton (Da).

Atomic weight

This is an obsolete term for relative atomic mass; its use is not recommended.

Atomisation

This is the process whereby a substance (usually in its standard state) is converted into gaseous atoms at infinite seaparation (which just means that they are far apart enough to not influence one another). In terms of energy this proces is normally quoted per mole of substance.

eg the enthalpy of atomisation of iron is the energy required to transform 1 mole of iron metal into 1 mole of gaseous iron atoms.

Fe_(s) --> Fe_(g)

Avogradro’s constant (L)

This should be defined as the number of particles in one mole of substance. The particles may be atoms, molecules, ions, electrons, etc. but should be clearly specified. Avogadro’s constant may be defined in terms of a particular number of particular number of particles, eg. 6.02 x 10²³. However, it should be noted that this value is continually being refined. The unit of Avogadro’s constant is mol^-1, and the term Avogadro’s number is to be discouraged because it implies a unitless quantity. Avogadro’s constant is also sometimes given the symbol, N_A.

Bar

A unit of pressure equal to 100 kPa and approximately equal to 750 mm Hg. It has been proposed that one bar be used as standard pressure, but this is not common practice.

Basic salt

A basic salt is a salt that retains a portion of the base from which it is derived. The retained part may be oxygen (eg. bismuth oxide chloride) or a hydroxyl group (eg. tin (II) hydroxide chloride. The formation of basic salts is an indication of weakness on the part of the base, and both the base and the basic salt are usually water-insoluble. Basic salts are often found in naturally occurring minerals, eg. malachite, basic copper (II) carbonate, Cu(OH)₂.CuCO₃.

Battery

A number of electrochemical cells connected in series. Often used in everyday speech to mean a single electrochemical cell.

Bi-

This prefix is used to describe

(1) an organic compound which contains two identical rings, eg. biphenyl. Sometimes replaced by the prefix di-;

(2) an acid salt, eg. sodium bicarbonate for sodium hydrogencarbonate. The use of bi-in this sense is not recommended;

(3) incorrectly, a species formed by condensation of two other species, eg. bichromate for dichromate. This term is no longer used.

Bond dissociation energy – see bond dissociation enthalpy

Bond dissociation enthalpy

The enthalpy increase that accompanies the homolytic fission of one mole of bonds in a covalently bonded species, to give individual atoms and/or radicals, with both the original species and the resulting fragments being in their standard states, and at standard temperature and pressure. The bond dissociation enthalpy refers to a specific bond in a specific compound. Note that the bond dissociation enthalpy is always positive. Also called (often/still) bond dissociation energy. The symbols E and D H_D are used for bond dissociation enthalpies.

Bond energy - see bond enthalpy

Bond enthalpy

The enthalpy increase that accompanies the homolytic fission of one mole of bonds in a covalently bonded species, to give individual atoms and/or radicals, with both the original species and the resulting fragments being in their standard states, and at stand temperature and pressure. The bond enthalpy is the average value of the bond dissociation enthalpy (see above) for the same kinds of bonds, but in a number of different compounds. Also called bond energy term or bond energy. The symbols E, E, D H_Dand D H_B are all used for bond enthalpy, which is sometimes confused with bond dissociation enthalpy.

Bond strength

A term that is used loosely to indicate bond enthalpies: the larger the bond enthalpy, the greater the strength of the bond. It is often used to compare different bonds qualitatively; thus, the intramolecular bonding in chlorine is stronger than the intermolecular bonding.

Buffer

An aqueous solution the pH of which remains nearly constant despite the addition of small amounts of acid or alkali. Buffers that maintain a pH of less than 7 are often made from a weak acid and one of its salts (eg. ethanoic acid and sodium ethanoate): buffers that maintain a pH of greater than 7 are often made from a weak alkali plus one of its salts (eg. ammonia solution and ammonium chloride). Acid salts anions, which are amphiprotic, any also give buffer solutions, eg. hydrogencarbonate ions that act as a buffer in blood. The use of strong acids as buffers at very low pH values, or strong alkalis at very high pH values, is often overlooked. A strong acid gives high concentration of hydrogen ions in solution, and this does not change appreciably when a small amount of base is added. Likewise a strong alkali contains a high concentration of hydroxide ions, which does not change appreciably when a small amount of acid is added.

Calorie

This is a term no longer used in the SI system to describe the quantity of heat required to raise the temperature of one gram of water by 1 ^oC (specifically from 14.5 to 15.5 ^ºC). One calorie = 4.148 J. The term finds everyday use in food, ie. the term food calorie or Calorie (1 Calorie = 1000 calories).

Carbocation

A species containing a trivalent carbon atom with one electron missing, it is thus associated with a positive charge eg. (CH₃)₃C+.

Carbonium ion - an alternative name for a carbocation

Cathode

The electrode at which reduction (ie. gain of electrons) takes place.

Cell

An apparatus that consists of a container with an electrolyte solution and electrodes immersed in it. The electrodes are connected into an external electrical circuit. There are two main types of cells : (1) galvanic cells, which produce electricity from chemical reactions, and (2) electrolytic cells, in which electrical energy is used to produce chemical reactions.

Celsius

A temperature scale in which the fixed points are the melting point of ice (0⁰C) and the boiling point of water (100 ⁰C) at standard pressure. One Celsius degree is equal to 1/100 the interval between the two fixed points, and is equal to a one degree interval on the absolute (Kelvin) temperature scale. The Celsius scale was formerly called Centigrade scale.

Centigrade – a commonly used alternative to Celsius

Charge centre

This term is used in Valence Shell Electron Pair Repulsion Theory. It commonly describes a single electron pair (bonding or non-bonding), or the two or three electron pairs found in double or triple bonds respectively, which are treated as a single centre for the purposes of the theory. It may also be used for a single electron in a half-filled orbital, or for bonds in delocalised systems.

Comproportionation

A type of redox reaction is which the same element in two different oxidation states react together to form one other, different, oxidation state, eg. iodide ions and iodate(V) ions react together to form iodine in the elemental state. (opposite disproportionation)

Concentration (symbol C, or indicated by square brackets around the substance under consideration, eg. [NaOH(aq)])

(1) A term used to describe the proportion of solute in a solution. It usually refers to the chemical amount of solute in a given volume of solution, however it may also refer to the mass of solute in a give volume of solution. Sometimes concentration is expressed as a percentage – meaning the mass of solute in 100 g of aqueous solution – or 100 cm³ if the solution is dilute. Occasionally the use of molal concentration is given : this refers to the chemical amount of solution in one kilogram of solvent. The term normality, referring to the concentration in terms of equivalent of solute per litre of solution, is obsolete.

The most common measure of concentration is mol dm^-3.

(2) The process of increasing the proportion of solute in a solution.

Condensation

This term is used to describe

(1) the process of changing from a gas to the liquid or solid state;

(2) a type of reaction in which two molecules combine to form a larger molecule, eliminating a small molecule as they do so, eg. carbonyl compounds react with hydrazine to form hydrazones, eliminating water as they do so.

Condensed formula

A formula often used to depict organic compounds, in which the structure is indicated, without fully displaying all the bonds, eg. the condensed formula for butan-2-ol is CH₃CH₂CH(OH)CH₃.

Conductivity

This term is used to describe

(1) the property of allowing electricity, heat, or sound to pass through a material. Also called conduction;

(2) a measure of the ease with which a substance (usually an electrolyte in solution) allows electricity to pass through it. Measured in units of W^-1 m^-1, it is the reciprocal of the resistivity. Formerly called the specific conductance;

(3) a measure of the ability of a substance to allow heat to flow through from a high temperature to a low temperature. Measured in units of J s^-1m^-1K^-1. Also called the thermal conductivity.

Conical flask - also called an Erlenmeyer flask. The term conical flask is recommended.

Contact process

The unit process in the manufacture of sulfuric acid in which oxygen and sulfur dioxide are reacted together over a vanadium pentoxide catalyst to form sulfur trioxide. Often incorrectly used to refer to the entire process of manufacturing sulfuric acid.

Coordination number

This term is determine as follows :
(1) in a metal the number of atoms that touch (or come closest to) any one particular atom. Thus close-packed structures have a coordination number of 12, but in body-centred cubic structures the coordination number is 8;

(2) in ionic lattices, the coordination number of a particular ion is the number of ions of the opposite charge that touch (or come closest to) it. Thus the coordination number of calcium ions is 8 but the coordination number of fluoride ions is 4 in the fluorite (CaF₂) structure;

(3) in complex ions the coordination number is determined by the number of coordinate bonds around the central ion. Thus the hexacyanoferrate (III) ion has a coordination number of 6. Note, however, that the coordination number is not necessary the same as the number of ligands around the central atom, as it is possible for some ligands to form more than one bond.

Coordinate bond

As alternative name for a dative covalent bond, often used in the context of complex ions.

Dative covalent bond

A chemical bond formed when two atoms share a pair of electrons, both of which may be considered to originate from the same atom.

Once formed a dative covalent bond cannot be distinguished from other covalent bonds in which the electrons in the bond are provided by different atoms. Also called a dative bond or a coordinate bond.

d- Block element

A metallic element with two s-electrons in its outer shell, and with between one and ten d-electrons in its penultimate shell. Sometime incorrectly used as a synonym for a transition metal: although all transition metals are d-block elements, the converse is not necessarily true.

Dehydration

This term is used to describe
(1) the process of removing water from a substance, for example, the removal of water of crystallisation from a hydrated salt;

(2) the process of removing the elements of water from a compound, ie. removing hydrogen atoms and oxygen atoms in the ratio 2 : 1, from a compound in which they are chemically bonded to other atoms, eg. ethanol is dehydrated to give ethene.

Delocalisation

A phenomenon in which valency electrons provided by individual atoms are no longer held in the near vicinity of that atom, but are mobile and shared by a number of atoms. Delocalisation occurs

(1) in metals, where electrons can move throughout the entire crystal structure;

(2) in organic compounds (and graphite) that have alternate double and single carbon-carbon bonds, orientated in such a way that p-electron can overlap, providing a pathway for electron movement;

(3) in certain inorganic species, such as nitrate and carbonate ions, where p-orbital overlap can occur. Delocalisation stabilises a structure, giving it a lower enthalpy than it would have if double and single bonds are arranged in such a way that orbital overlap cannot occur.

Diacidic

Describes a base one mole of which is capable of neutralizing two moles of protons, eg. calcium hydroxide may be described as diacidic. The term diacid is sometimes also used in the same context, but this should be avoided as it is also occasionally used to describe a diprotic acid, or (more often) a molecule containing two acid groups such as ethanedioic acid.

Diaphragm cell - see membrane cell

Dibasic

Describes an acid one mole of which is capable of producing two moles of protons in a neutralisation reaction. Also called diprotic.

Dipole-dipole forces

These forces are a type of intermolecular bonding caused by attractions between permanent dipoles in polar molecules. Also known as dipole forces or dipole-dipole interactions or permanent dipole-dipole interactions.

Diprotic - see dibasic

Discharge

This term is used to describe

(1) the discharge seen at an electrode during electrolysis, eg. chlorine is discharged at the anode during electrolysis of molten sodium chloride;

(2) the electrical discharge in a discharge tube;

(3) in common language, the removal of charge, ie. send away.

Dispersion forces

Forces between temporary dipoles induced in atoms or molecules also known as London forces or temporary dipole-dipole interactions. See van der Waal’s forces.

Disproportionation

A type of redox reaction in which an element in one particular oxidation state is simultaneously oxidised and reduced to give two products in different oxidation states, eg. chlorine molecules react with water to form chloride ions and chlorate (I) ions.

Dissociation

This term is used to describe

(1) separation of a compound into ions when it is dissolved in a polar solvent. For example molecules of sulfuric acid are dissociated into hydrogen ions and sulfate ions in water. Ionic compounds, such as salts, may also dissociate if the ion-dipole interactions are sufficiently strong. Dissociation is often reversible and equilibrium conditions prevail;

(2) a reversible process in which a compound splits up into smaller species. Thus ammonium chloride is dissociated into ammonia and hydrogen chloride when heated, but the ammonium chloride is formed again on cooling.

Dissociation constant - see acid ionisation constant

Double decomposition - see double displacement

Double displacement

An exchange reaction between two ionic compounds in solution, in which the cations and anions change partners, thus :

AB + CD º AD + CB The products of the reaction may be insoluble or volatile compounds, so that they are removed from solution, allowing the equilibrium to swing in favour of the products. Not that the products are not necessarily ionic, thus water is formed by the reaction of hydrogen and hydroxide ions in the neutralisation reaction. Double displacement is also called double decomposition and metathesis. Also see ionic precipitation.

Efflorescent

This term is used to describe :
(1) a substance that loses water of crystallisation, becoming powdery in the process. Thus hydrated sodium carbonate is efflorescent, gradually losing water of crystallisation to the atmosphere to become anhydrous sodium carbonate;

(2) a substance that becomes encrusted with powder or crystals as a result of chemical change or the evaporation of a solution. For example, copper sulfate solution may gradually evaporate to give an encrustation that may climb up and over the sides of the container.

Electrochemical cell

This is often defined as a cell in which a spontaneous chemical reaction is used to produce electrical energy, and may also be called a galvanic cell or a voltaic cell. The term electrochemical cell is sometimes used for electrolysis cells as well, so it is recommended that the term galvanic cell is used for a cell in which a chemical reaction produces electrical energy.

Electrochemical series

A list of elements written in order of their standard reduction potentials, having those element with the most negative reduction potentials written first. As this list refers to reductions in aqueous solutions, it is slightly different from the reactivity series which is based on the reduction of metal oxides. The electrochemical series may also be extended to include non-metals and other reduction systems. The electrochemical series is sometimes written with the most positive reduction potentials first, or (in very old-fashioned books) as oxidation potentials. However, the use of such lists is not recommended.

Electron affinity

(1) The first electron affinity of an element is the enthalpy change that occurs when one electron is gained by each atom in a mole of gaseous atoms of the element to give one mole of ions, each with a single negative charge, at standard temperature and pressure.

(2) This term is also used to describe the enthalpy increases that occur when subsequent electrons are gained. For example the second electron affinity would refer to the gain of one electron by each ion, each with a single negative charge, in a mole of such ions in the gas state, to give one mole of ions with a double negative charge in the gas state. Note that according to these definitions most common first electron affinities are negative, but second electron affinities are often positive.

Electron dot formula - see Lewis structure

Electronegativity

The ability of an atom to attract electrons towards itself in a diatomic bond. There are two different Electronegativity scales:

(1) Mulliken’s scale in which the Electronegativity of an atom is the arithmetic mean between the ionisation energy and the electron affinity;

(2) the more commonly used Pauling scale, in which all values are measured relative to fluorine, which has the maximum Electronegativity of 4.0.

Using a Pauling scale, it is possible to estimate the degree of ionic character of a covalent bond, by calculating the difference in Electronegativity of the two atoms in the bond: a difference of $ 1.8 indicates that the bond is essentially ionic.

Element

This term is used to describe
(1) on a macro scale, the simplest form of substance: it cannot be further broke down by chemical means;

(2) on an atomic scale, a substance the atoms of which all have the same nuclear charge. All atoms of the same element display the same chemical characteristics.

Empirical formula - also called simplest formula

End point - see equivalence point

Enthalpy change of atomisation

Note that different definitions apply according to whether the substance being atomised is an element or a compound. However, the enthalpy changes involved in either case are positive as they involve bond breaking reactions.

(1) The standard enthalpy change of atomisation of an element, symbol D Hq_atm, is the enthalpy increase that takes place when one mole of gaseous atoms is made from the element in the defined physical state under standard conditions.

(2) The standard enthalpy change of atomisation of a compound, symbol D Hq_atm,is the enthalpy increase that takes place when one mole of the compound in the defined physical state, is broken down into gas atoms under standard conditions.

Equilibrium expression - see equilibrium law

Equilibrium law

The equilibrium law for a reaction is the expression in which the equilibrium constant is equal to a fraction in which the numerator is the product of the concentrations of the substances on the right of the equation, each raised to a power equal to its coefficient in the chemical equation. The denominator is the product of the concentrations of the substances on the left of the equation, each raised to a power equal to its coefficient in the chemical equation. The equilibrium law is also called the equilibrium expression.

Equivalence point

That point in a titration where stoichiometrically equivalent amounts of reactants have been added.

Erienmeyer flask - see conical flask

Ether

(1) The common name for ethoxyethane.

(2) Any organic compound containing a C-O-C linkage.

Fission

The term is used to describe

(1) the breaking of a chemical bond, eg. Cl₂(g) ® 2Cl^. (g);

(2) the breaking up of a nucleus.

Fusion

The term is used to describe

(1) melting, for example, in enthalpy change of fusion

(2) the combining of two nuclei, eg. ²₁D + ²₁D ®⁴₂He

Galvanic cell

An electrochemical cell that generates electricity by means of spontaneous redox reaction, also known as a Voltaic cell.

Giant molecular solids – see network covalent solids

Heats of reaction – see enthalpy changes

Hess’ law

This is also called the law of constant heat summation/the first law of thermodynamics

Hydrolysis

The reaction of a substance with water to form two or more products,

eg. CH₃COCl + H₂O ® CH₃CO₂H + HCl;

Hydronium ion

This ion, H₃O⁺, is also oxonium ion and hydroxonium ion.

Hydroxonium ion - see hydronium ion

Hydroxyl

Describes a group having a hydrogen and oxygen atom covalently bonded together. The hydroxyl group is often covalently bonded via the oxygen atom into a molecule or ion, as in an alkanol or a hydrogensulfate ion. However, hydroxyl is also often used as a synonym for hydroxide.

Ionic precipitation

A form of double decomposition where ions combine to produce a precipitate.

Ionisation

This term is used to describe

(1) the formation of ions, eg. Ar ® Ar⁺ + e;

(2) the dissolving of a solid in water to form ions eg. CH₃CO₂H CH₃CO₂^-(aq) + H⁺(aq)

Ionisation constant – see acid ionisation constant

Kelvin Scale

A temperature scale with intervals measured in Kelvin. Also known as the absolute temperature scale.

Kinetics

A study of the rates of chemical reactions. This is separate from the kinetic theory.

The kinetic theory

The theory in which the behaviour of matter is explained in terms of the movement of small particles. This is separate from kinetics.

Lattice enthalpy

The enthalpy change that occurs when one mole of a solid ionic crystal is separated into each of its component ions in the gaseous state, at standard temperature and pressure. Lattice enthalpy defined in this way will always have values that are positive. Many texts define it the opposite way, ie. for the change from separate ions to ionic crystal. Its value is then negative.

Law of conservation of energy – see Hess’ law

Lewis structure

A diagram showing the covalent bonds in a molecule or ion by using the symbol(s) of the element(s) involved and some representation of the valency electrons. This representation can be by dots, crosses, a combination of dots and crosses or by using a line to represent a pair of electrons.

eg. H : H, H^x₀ H, or H - H

Also known as a Lewis formula or electron dot formula.

Litre (symbol L, dm³)

A non-SI unit of volume equivalent to one dm³, which is often still used to indicate the capacity of containers. (The definition of the litre as 1000.028 cm³ was abandoned in 1964). It is not recommended for use in calculations as it is not properly coherent with SI units.

London forces - see dispersion forces

Macromolecular solids - see network covalent solids for recommended term.

Mass of one mole - see molar mass

Membrane cell

The most modern process used for the electrolysis of sodium chloride solution in the chlor-alkali industry. It replaces the diaphragm cell.

Molar

This means ‘divided by the amount of substance in moles’. Thus the molar volume is the volume per mole of substance. There are some exceptions to this definition. For example, molar conductivity where molar means ‘divided by the concentration in mol dm^-3". It is recommended that the term molar is not used to refer to solutions, where it is often incorrectly used to indicate the chemical amount in a given volume of solution.

Molar mass, M

This is the mass per mole of substance, and has the units of g mol^-1. The molar mass should always be accompanied by a statement indicating the nature of the particles in the substance, or the formula of the substance, eg. for chlorine, it should be specified whether the molar mass refers to chlorine atoms, Cl, or chlorine molecules Cl₂, as the molar masses are different for the two species.

Molar volume, V_n

This is the volume per mole of substance, and usually has the units of dm³.mol^-1. The conditions under which the molar volume is measured should always be stated, and this is particularly important for gases, where small changes in temperature and pressure can make a significant difference to the volume. The molar volume of an ideal gas is 22.4 dm³ mol^-1 at 273 K and 1.00 atm, which converts to approximately 24 dm³ mol^-1 at 298 K and 1.00 atm.

Molarity

This refers to the concentration of a solution in mol dm^-3. The term is now considered to be obsolete. The abbreviation M for the units mol dm^-3 is still used, though it is not recommended.

Network covalent solids

Atoms bonded together covalently throughout the solid, eg. diamond, silicon oxide. Also called giant molecules, macromolecular solids.

Normal temperature

At one time this was taken as 18 ^oC. However it has also been used to mean standard temperature (0^oC) or stand ambient temperature (25 ^oC). Nowadays it is sometimes used loosely to mean room temperature. Because of these different connotation, it is recommended that the term be avoided.

Oxonium ion - see hydronium ion

Periodic Table

The table in which the elements are arranged in groups and periods. Also sometimes called the Periodic Classification.

Periodic Classification - see Periodic Table

An abbreviation commonly used to indicate an alkyl or aryl group in formulae of organic compounds. Also used for the gas constant.

Rate constant

This term is sometimes called the velocity constant.

Rate equation - see rate expression

Rate expression

An equation relating the rate of the reaction to the rate (or velocity) constant, and the concentrations of the reactants, also known as the rate law or rate equation.

Rate law - see rate expression

Reactivity series

A series of chemical elements (usually metals and hydrogen) arranged in order of their tendency to lose electrons, ie. their strength as reducing agents, also known as the activity series (or in the past electromotive series – this term is no longer used). See also the closely related electrochemical series.

Resonance hybrids

Alternative methods of representing the bonding in molecules or ions for a given arrangement of atoms. If the various possible arrangements differ only a littler in energy, than the actual bonding is a mixture of the various possible representations, and will have a lower energy than any of the individual forms. Also called canonical forms.

Room temperature and pressure - see standard ambient temperature and pressure

Saturated

This is a term used to describe

(1) organic compounds that contain no multiple bonds (eg. C=C, C=O, C/ N);
(2) solutions of a solute and solvent in which, at a given temperature, no more solute will dissolve.

Simplest formula - see empirical formula

Standard ambient temperature and pressure

Standard ambient temperature is 298 K (25.0^oC) and standard pressure is one atmosphere (101 325 Pa). These are the standards most commonly used for chemical reactions. Compare with standard temperature and pressure. This is also called room temperature and pressure, and is the standard commonly used in thermodynamics.

Standard electrode potentials

The electromotive force of a half-cell connected to a standard hydrogen electrode, measured under standard conditions. The sign of the electrode potential being measured is positive if reduction occurs at it, but negative if the reduction occurs at the hydrogen electrode. Also called the standard reduction potential or the standard redox potential.

Standard pressure

This is usually taken as one standard atmosphere or 101 325 Pa (101.325 kPa). The use of a pressure of one bar (100 kPa) has also been proposed, but this is not recommended.

Standard redox potential - see standard electrode potential

Standard reduction potential - see standard electrode potential

Standard temperature and pressure

Standard temperature is often taken as 273.15 K (0.0 ^oC) and standard pressure as one atmosphere (101 325 Pa). This is the standard commonly used in gas calculations. Compare with standard ambient temperature and pressure.

Transition metal

An element the atoms of which have an incomplete set of d-electrons in their penultimate shell in one or more of their oxidation states. This includes all elements in the d-block except those with complete d-orbitals, such as zinc. Scandium is sometimes excluded from the transition metals because its ions have empty d-orbitals, and thus do not exhibit transition characteristics.

Transition state

An unstable arrangement of atoms that exists for only a moment in the course of a chemical reaction.

Universal gas law

Also known as the ideal gas equation.

Unsaturated

This term is used to describe

(1) substances that contain multiple bonds, eg. C=C, C=O, C=N. The term is used most commonly in association with organic compounds;

(2) solutions of a solute and a solvent, at a given temperature, into which more solute can be dissolved.

Van Der Waals’ forces

These are weak attractive forces between molecules (or intermolecular bonds). The term is sometimes used to refer to all types of intermolecular forces, ie. dispersion forces, dipole-dipole interactions and hydrogen bonds, or just the first of these, or just the first two! The term van der Waals’ forces is mostly restricted to the first, ie. its recommended use is for dispersion forces.

Velocity constant - see rate constant

VSEPR theory

Abbreviation for valence shell electron pair repulsion theory.