F325 Flashcards
Acid dissociation constant
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Bond dissociation enthalpy
The enthalpy change that takes place when breaking by homolytic fission 1 mol of a given bond in the molecules of a gaseous species.
Acid–base pair
A species that is a proton, H+, donor. A pair of two species that transform into each other by gain or loss of a proton.
Brønsted–Lowry acid
A species that is a proton, H+, donor.
Activation energy
A species that is a proton, H+, acceptor. The minimum energy required to start a reaction by the breaking of bonds.
Brønsted–Lowry base
A species that is a proton, H+, acceptor.
Adsorption
The process that occurs when a gas, liquid orsolute is held to the surface of a solid or, more rarely, aliquid.
Buffer solution
A system that minimises pH changes on addition of small amounts of an acid or a base.
Alkali
A type of base that dissolves in water to form hydroxide ions, OH- (aq) ions.
Catalyst
A substance that increases the rate of a chemical reaction without being used up in the process.
Average bond enthalpy
The average enthalpy change that takes place when breaking by homolytic fission 1 mol of a given type of bond in the molecules of a gaseous species.
Complex ion
A transition metal ion bonded to one or more ligands by coordinate bonds (dative covalent bonds).
Boltzmann distribution
A diagram showing the distribution of energies of the molecules at a particular temperature.
Conjugate acid
A species formed when a proton is added to a base.
Conjugate base
A species formed when a proton is
added to an acid.
Electron shielding
The repulsion between electrons in different inner shells. Shielding reduces the net attractive force from the positive nucleus on the outer shell electrons.
Coordination number
The total number of coordinate bonds formed between the central metal ion and any ligands.
Electronegativity
A measure of the attraction of a bonded atom for the pair of electrons in a covalent bond.
Displacement reaction
A reaction in which a more reactive element displaces a less reactive element from an aqueous solution of its ions.
End point
The point in a titration at which there are equal concentrations of the weak acid and conjugate base forms of the indicator. The colour at the end point is midway between the colours of the acid and conjugate base forms.
Disproportionation
The oxidation and reduction of the same species in a redox reaction.
Endothermic
A reaction in which the enthalpy of the products is greater than the enthalpy of the reactants, resulting in heat being taken in from the surrounding.
Dynamic equilibrium
The equilibrium that exists in a closed system when the rate of the forward reaction is equal to the rate of the reverse reaction.
(Standard) enthalpy change
of atomisation
The enthalpy change that takes place when one mole of gaseous atoms forms from the element in its standard state.
(First) electron affinity
The enthalpy change required to add one electron to each atom in one mole of gaseous atoms to form one mole of gaseous 1– ions.
(Standard) enthalpy change of
combustion
The enthalpy change that takes place when one mole of a substance reacts completely with oxygen under standard conditions, all reactants and products being in their standard states.
(Second) electron affinity
The enthalpy change required to add one electron to each ion in one mole of gaseous 1– ions to form one mole of gaseous 2– ions.
(Standard) enthalpy change of formation
The enthalpy change that takes place when one mole of a compound is formed from its constituent elements in their standard states under standard conditions.
(Standard) enthalpy change of hydration
The enthalpy change that takes place when one mole of isolated gaseous ions is dissolved in water, forming one mole of aqueous ions, under standard conditions.
Entropy, S
The quantitative measure of the degree of disorder in a system.
(Standard) enthalpy change of neutralisation
The energy change that accompanies the neutralisation of an aqueous acid by an aqueous base to form one mole of H2O(l), under standard conditions.
(Standard) entropy change of
reaction
The entropy change that accompanies a reaction in the molar quantities expressed in a chemical equation under standard conditions, all reactants and products being in their standard states.
(Standard) enthalpy change of reaction
The enthalpy change that accompanies a reaction in the molar quantities expressed in a chemical equation under standard conditions, all reactants and products being in their standard states.
Equilibrium law
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(Standard) enthalpy change of solution
The enthalpy change that takes place when one mole of a compound is completely dissolved in water under standard conditions.
Equivalence point
The point in a titration at which the volume of one solution has reacted exactly with the volume of the second solution. This matches the stoichiometry of the reaction that is taking place.
Enthalpy cycle
A diagram showing alternative routes between reactants and products that allows the indirect determination of an enthalpy change from other known enthalpy changes using Hess’s law.
Exothermic
A reaction in which the enthalpy of the products is smaller than the enthalpy of the reactants, resulting in heat loss to the surroundings.
Enthalpy profile diagram
A diagram for a reaction to compares the enthalpy of the reactants with the enthalpy of the products.
Free energy change
The balance between enthalpy, entropy
and temperature for a process:
ΔG = ΔH – TΔS
A process can take place spontaneously
when ΔG
Enthalpy, H
The heat content that is stored in a chemical system.
Half-life
The time taken for the concentration of a reactant to reduce by half.
Hess’s Law
If a reaction can take place by more than one route and the initial and final conditions are the same, the total enthalpy change is the same for each route.
Ionic product of water, Kw
Kw = [H+(aq)] [OH–(aq)]
At 25°C, Kw = 1.00 × 10–14 mol2dm–6.
Heterogeneous catalysis
A reaction in which the catalyst has a different physical state from the reactants; frequently reactants are gases whilst the catalyst is a solid.
(First) ionisation energy
The energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.
Heterogeneous equilibrium
An equilibrium in which the species making up the reactants and products are in different physical states.
(Second) ionisation energy
The energy required to remove one electron from each ion in one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions.
Homogeneous catalysis
A reaction in which the catalyst and reactants are in the same physical state, which is most frequently the aqueous or gaseous state.
Lattice enthalpy
The enthalpy change that accompanies the formation of one mole of an ionic compound from its gaseous ions under standard conditions.
Homogeneous equilibrium
An equilibrium in which all the species making up the reactants and products are in the same physical state.
Le Chatelier’s Principle
When a system in dynamic equilibrium is subjected to a change, the system readjusts itself to minimise the effect of the change and to restore equilibrium.
Initial rate of reaction
The change in concentration of a reactant or product per unit time at the start of the reaction: t = 0.
Ligand
A molecule or ion that can donate a pair of electrons to a transition metal ion.
Intermediate
A species formed in one step and used up in a subsequent step and so never seen as either a reactant or a product.
Ligand substitution
A reaction in which one ligand in a complex ion is replaced by another ligand.
Limiting reagent
The substance in a chemical reaction that runs out first.
Periodicity
A regular periodic variation of properties of elements with atomic number and position in the Periodic Table.
Order
The power to which the concentration of the reactant is raised in the rate equation.
pH
pH = –log[H+(aq)]
[H+(aq)] = 10–pH.
Nucleophile
An atom (or group of atoms) which is attracted to an electron-deficient centre or atom, where it donates a pair of electrons to form a new covalent bond.
Rate constant, k
The constant that links the rate of reaction with the concentrations of the reactants raised to the powers of their orders in the rate equation.
Overall order
The sum of the individual orders: m + n.
Rate equation
For a reaction: A + B → C, the rate equation is:
rate = k[A]m[B]n
m is the order of reaction with respect to A.
n is the order of reaction with respect to B.
m + n = overall order.
Oxidation
Loss of electrons or an increase in oxidation number.
Rate of reaction
The change in concentration of a reactant or a product in a given time.
Oxidation number
A measure of the number of electrons that an atom uses to bond with atoms of another element. Oxidation numbers are derived from a set of rules.
Rate-determining step
The slowest step in the reaction mechanism of a multi-step reaction.
Oxidising agent
A reagent that oxidises (takes electrons from) another species.
Reaction mechanism
A series of steps that, together, make up the overall reaction.
Redox reaction
A reaction in which both reduction and oxidation take place.
Reducing agent
A reagent that reduces (adds electrons to) another species.
Strong acid
An acid that completely dissociates in solution.
Reduction
Gain of electrons or a decrease in oxidation number.
Thermal decomposition
The breaking up of a chemical substance with heat into at least two chemical substances.
Specific heat capacity, c
The energy required to raise the temperature of 1g of a substance by 1°C.
Transition element
A d-block element which forms an ion with an incomplete d sub-shell.
Stability constant, Kstab
The equilibrium constant for an equilibrium existing between a transition metal ion surrounded by water ligands and the complex formed when the same ion has undergone a ligand substitution reaction.
Weak acid
An acid that partially dissociates in solution.
Standard electrode potential
The e.m.f. of a half cell compared with a standard hydrogen half cell, measured at 298K with solution concentrations of 1 moldm–3 and a gas pressure of 101 kPa (1 atmosphere).
