Chemical Definitions: December 2008

(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.

Chemical Definitions

Wednesday, December 3, 2008

Chemical Physics Definitions

Followers

Blog Archive

About Me