Definitions- define each term given Flashcards
Atomic number
The number of protons in the nucleus of an atom.
Mass number
The number of protons and neutrons in the nucleus of an atom.
Isotopes
Atoms of the same element so the same atomic number with different numbers of neutrons so different mass numbers.
Relative atomic mass, Ar
The ratio of the average mass of one atom of an element to one-twelfth of the mass of an atom of carbon-12.
Relative molecular mass, Mr
The ratio of the average mass of one molecule of a compound to one-twelfth of the mass of an atom of carbon-12.
Mole
The amount of any substance containing as many particles as there are carbon atoms in exactly 12g of the carbon-12 isotope.
Avogadro’s constant, Na
The number of particles per mole (6.02 x 10^23).
Empirical formula
The simplest whole number ratio of atoms of each element present in a compound.
First ionisation energy
The energy required to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.
Successive ionisation energy
The energy required to remove each electron in turn.
Electronegativity
The measure of the pull of an atom on a pair of electrons in a chemical bond.
OR
The ability of an atom to pull electron density towards itself within a covalent bond.
Ionic bond
A bond formed by the electrostatic forces of attraction between a positively charged ion and a negatively charged ion.
Covalent bond
Formed when atoms share a pair of electrons.
Metallic bonding
Bonding between atoms in a metal crystal. Each metal atom contributes electrons from its outer shell to a sea of delocalised electrons.
Dative covalent (coordinate) bond
A bond formed when one atom contributes both electrons in a covalent bond
Acid
A proton donor.
Anion
A negatively charged ion.
Period
A horizontal row of elements in the periodic table.
Salt
Compound formed when H+ from an acid is replaced by a metal ion or other positive ion.
Disproportionation
The oxidation and reduction of the same element in a redox reaction.
Thermal decomposition
The breaking up of a chemical substance with heat into at least two chemical substances.
Oxidation
The loss of electrons or increase in oxidation number.
Cation
A positively charged ion.
Molecular formula
The actual number of atoms of each element in a molecule.
Enthalpy change (🔼H)
The heat energy change in a reaction under constant pressure.
Standard conditions
100KPa (about 1 atm) pressure and a stated temperature, usually 298K.
Standard enthalpy of combustion
The enthalpy change of a compound when 1 mole of the compound is burned completely in oxygen under standard conditions, with all reactants and products in their standard state.
Standard enthalpy of formation
The enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions.
Relative isotopic mass
The weight mean mass of an atom of an element compared with 1/12 of the mass of carbon 12
Relative formula mass
The formula mass is relative to the mass of an atom of carbon 12 which is defined as exactly 12.
Molar mass
The number of atoms, molecules, ions or other entities in one mole of a substance
Anhydrous
Anhydrous means ‘without water’
Hydrated
Contains water
Water of crystallisation
Water molecules which make up part of the crystal structure of a compound
Base
A proton acceptor
Alkali
A soluble base
Oxidation state
The state of oxidation or reduction shown by an element in its chemistry
Electron orbital
A subdivision of a sub shell. Each orbital contains upto 2 electrons
Electron shell
An energy level representing the distance of a group of electrons from the nucleus of an atom.
Permanent dipole
An intermolecular force between two polar molecules
Polar covalent bond
When two atoms with different electronegativities bond covalently
Intermolecular force
Attractive forces between molecules
Hydrogen bond
A type of attraction between molecules which is much stronger than all the types of intermolecular forces but much weaker than covalent bonding
Periodicity
A repeating pattern of properties shown across different periods
Element
All the atoms of an element have the same number of protons in the nucleus.
Reduction
A gain of electrons
Reducing agent
Something that donates electrons
Polymerisation
The joining together of monomers to form long chains.
Electrophile
Electron pair acceptor
Addition
Reaction which increases number of substituents
Dehydration
The elimination of water from a compound
Structural isomers
Compounds with the same molecular formula but different structural formulae
Positional isomers
Compounds with the same molecular formula but different structures due to different positions of the same functional group on the same carbon skeleton
Biofuel
A fuel made from plants or organic matter
Oxidation state
The charge on the ion or element or atom
Oxidising agent
A substance which accepts electrons
Dynamic equilibrium
Rate of the forward reaction equals rate of the backwards reaction and concentrations of reactants and products remain constant
Compromise temperature
Balance between rate and yield
Activation energy
Minimum energy to start a reaction
Catalyst
Speeds up a reaction but is chemically unchanged at the end
Mean bond enthalpy
Enthalpy to break a covalent bond varies between compounds so an average value is used
Stereoisomers
Compounds with the same structural formula but a different arrangement of atoms in space
Rate of reaction
The change in concentration per unit of time
Nucleophile
An electron pair donor
Carbon neutral
An activity which has no net carbon emissions to the atmosphere
Hess’s Law
Enthalpy change is independent of the route taken
Racemic mixture
Equal mixture of enantiomers
Bronsted-Lowry Acid
A proton donor
Weak acid
A species which partially dissociates in solution to give H+ ions
Strong acid
A species which fully dissociates in solution to give H+ ions
Bronsted-Lowry Base
A proton acceptor
Strong base
A species which fully dissociates in solution to give OH- ions
Weak base
A species which partially dissociates in solution to give OH- ions
pH
The negative logarithm to the base 10 of the hydrogen ion concentration
Kw
The product the concentrations of hydrogen and hydroxide ions a solution
Buffer
A solution which can resist a large pH change on addition of small amounts of acid or alkali
Chiral
An atom attached to four different groups (or a molecule containing an atom attached to four different groups)
Enantiomers
Molecules which are non-superimposable mirror images of each other
Optical Activity
The ability to rotate the plane of plane-polarised light
Second ionisation enthalpy
Enthalpy change when one electron is removed from each of a mole of free unipositive ions of that element in the gaseous state
Enthalpy of Atomisation of an element
Enthalpy change when one mole of free gaseous atoms is produced from that element in its standard state.
Enthalpy of Atomisation of a Compound
Enthalpy change when one mole of a compound in its standard state is converted into free gaseous atoms
First Electron Affinity
Enthalpy change when one electron is added to each of a mole of free gaseous atoms of that element
Second electron affinity
Enthalpy change when one electron is added to each of a mole of free gaseous uninegative ions of that element.
Bond Dissociation Enthalpy
Mean Enthalpy change when one mole of covalent bonds is broken homolytically, resulting in free gaseous atoms
Lattice Enthalpy
Enthalpy change when one mole of an ionic compound is formed from its free gaseous ions
Lattice dissociation Enthalpy
Enthalpy change when one mole of an ionic compound is completely dissociated into free gaseous ions.
Enthalpy of Hydration
Enthalpy change when one mole of free gaseous ions is added to an excess of water.
Enthalpy of solution
Enthalpy change when one mole of an ionic compound is completely dissolved in an excess of water
Entropy
A measure of the amount of disorder in a substance
Spontaneous Reaction
A reaction for which the free energy change is negative.
Amphoteric
A substance which is able to react with acid and alkalis
Oxidation number
The charge that would exist on an atom if all the bonding were completely ionic
Standard electrode potential
The emf of a cell in which the left-hand electrode is the standard hydrogen electrode and the right-hand electrode is the standard electrode in question.
Ligand
A species which can use its lone pair of electrons to form a co-ordinate bond with a metal ion.
Co-ordination number
The total number of co-ordinate bonds formed between the metal ion and the ligands in a complex
Complex
A species containing a metal ion attached to one or more ligands by means of coordinate bonds
Bidentate ligand
A ligand which uses two lone pairs of electrons to from two coordinate bonds with a metal ion.
Multidentate ligand
A ligand which uses more than two lone pairs of electrons to form more than two coordinate bonds with a metal ion
Homogeneous Catalyst
A catalyst in the same physical state as the reactants.
Heterogeneous Catalyst
A catalyst in a different physical state to the reactants.
Lewis Acid
An electron pair acceptor
Lewis base
An electron pair donor
