AQA A-Level Chemistry Paper 1 Definitions

Relative Atomic Mass

the average mass of one atom of an element divided by 1/12th of the mass of 1 atom of carbon 12

Mass Number

number of protons and neutrons in an atom's nucleus

Atomic Number

number of protons in an atom's nucleus

Isotope

atoms with the same number of protons but different number of neutrons

Electrospray Ionisation

method of ionisation in mass spectrometry. dissolve sample in polar volatile solvent. sample gains a proton from the solvent. injected into spectrometer using a high voltage needle to create a fine spray. solvent evaporates = XH+ ion

X add H+ = XH+

Electron Gun (Electron Impact) Ionisation

method of ionisation in mass spectrometry. fire a high energy electron at the sample. knocks off an electron to form a positive ion. MUST BE GAS PHASE.

X(g) add e- = X+ add 2e-

Isoelectronic

ions with the same electronic structure/ compounds with the same number of electrons

1st Ionisation Energy

energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous ions with singly positive charges

Empirical Formula

the simplest whole number ratio of atoms of each element present in a compound

% Atom Economy

% of the mass of atoms of reactants converted into useful products

Ionic Bond

electrostatic force of attraction between oppositely charged ions

Ionic Lattice

the regular repeating pattern of oppositely charged ions next to each other

Covalent Bonding

a shared pair of electrons between non metals

Coordinate (Dative Covalent) Bond

a covalent bond where the pair of electrons in the bond come from a single atom (denoted by an arrow)

Lone Pair

pair of electrons not involved in bonding

Metallic Bonding

the attraction between delocalised electrons and positive ions arranged in a crystalline structure or lattice

Electronegativity

the power of an atom to attract a pair of electrons in a covalent bond

Polar Bond

when two atoms in a covalent bond have different electronegativities and the two electrons are not shared equally. results in delta + and delta -. the dipole moments do not cancel

Van Der Vaals Forces

induced dipole - induced dipole forces. occur due to electrons continually moving around so there is an uneven distribution of electrons. causes temporary dipole which induces another temporary dipole within a molecule.

Permanent Dipole - Permanent Dipole Forces

permanent attractions between polar molecules

Hydrogen Bonding

the strongest intermolecular force. hydrogen attached to a very electronegative atom with a lone pair. linear, lone pairs, deltas required.

Standard Enthalpy Change

heat energy change at constant pressure under standard conditions

Standard Conditions

standard states, 100kPa, 298K

Endothermic

heat energy is absorbed from the surroundings (gets cold) +

Exothermic

heat energy is released to the surroundings (gets hot) -

Activation Energy

the minimum energy required for a chemical reaction to occur (bonds to break)

Hess' Law

the enthalpy change of a chemical reaction is independent of the route taken

Standard Enthalpy of Formation

the enthalpy change when 1 mole of a substance is formed from its elements under standard conditions, all substances in standard states

Standard Enthalpy of Combustion

the enthalpy change when 1 mole of a substance completely combusts in excess oxygen under standard conditions, all substances in standard states

Bond Enthalpy

enthalpy change to break 1 mole of covalent bonds in gaseous state

Mean Bond Enthalpy

mean bond enthalpy across a range of compounds

Collision Theory

for a reaction to take place particles must collide with enough activation energy

Oxidation

loss of electrons

Oxidising Agent

electron acceptor

Reduction

gain of electrons

Reducing Agent

electron donor

Disproportionation

when one species is simultaneously oxidised and reduced

Redox

when oxidation and reduction take place in the same reaction

Oxidation State

the condition of an atom expressed by the number of electrons that the atom needs to reach its elemental form

Half Equation

show the movement of electrons in redox reactions

First Electron Affinity

enthalpy change when each atom in one mole of gaseous atoms gains one electron to form one mole of gaseous 1- ions

Second Electron Affinity

enthalpy change when each ion in one mole of gaseous 1- ions gains one electron to form one mole of gaseous 2- ions

Enthalpy of Atomisation

enthalpy change when 1 mole of gaseous atoms is produced from an element in its standard states

Enthalpy of Hydration

enthalpy change when 1 mole of gaseous ions are dissolved in water

Enthalpy of Solution

enthalpy change when one mole of an ionic solid dissolves in an amount of water large enough so that the dissolved ions are well separated and do not interact with each other

Bond Dissociation Enthalpy

enthalpy change when one mole of covalent bonds is broken in the gaseous state

Lattice Enthalpy of Formation

enthalpy change when 1 mole of a solid ionic compound is formed from its gaseous ions under standard conditions

Lattice Enthalpy of Dissociation

enthalpy change when one mole of a solid ionic compound is broken up into its constituent ions in the gas phase

Covalent Character

a compound has covalent character when the negative ion's electron cloud is distorted because it has been polarised by the positive ion. usually with small highly charged positive ions and large highly charged negative ions

Perfect Ionic Model

assumption that the ions involved in a compound are perfect spheres. big difference between theoretical and calculated value = covalent character

Entropy

a measure of disorder. positive entropy change means the reaction is favourable

Gibbs Free Energy

the energy of a system that is available to do work at a constant temperature and pressure. helps to decide if a reaction is feasible. if gibbs free energy change is less than or equal to zero = feasible

Dynamic Equilibrium

the rate of the forward and backward reactions are equal and the concentrations of reactants and products are constant

Le Chatelier's Principle

the position of the equilibrium in a dynamic equilibrium moves to oppose the change

Kc for Homogeneous Systems

equilibrium constant where everything is in the same phase

Kp for Homogeneous Systems

the equilibrium constant for a homogenous system where everything is in the gas phase

Partial Pressure

the pressure a gas would exert if it were alone in the container

Order of a Reaction

how the overall rate of reaction is affects by the concentration of that reactant

Zero Order

the reaction rate is unchanged by changing the concentration of A

First Order

the reaction rate is proportional to the chnage in concentration of A

Second order

the reaction rate is proportional to the change in concentration of A squared

Rate Determining Step

the slowest step of a mechanism of a reaction. species after the RDS do not appear in the rate equation

Overall Order

the sum of the orders of all reactants in a chemical reaction

Half Cell

metal immersed in a solution of its own ions

Electrochemical Cell

combining two half cells together, connected by a salt bridge and a high resistance voltmeter

SHE

standard hydrogen electrode known as primary standard, used to compare all other potentials. must be standard conditions of 1 moldm-1, 298K, 100kPa

hydrogen gas, platinum electrode, hcl

2H+ add 2e- = H2

EMF

electromotive force (ER-EL)

E nought

standard potential

Salt Bridge and Cell Notation

a tube or solution soaked filter paper that completes the circuit and allows the transfer of ions.

no solid metal in the cell notation? add platinum electrode

, = same phase

| = phase boundary

RO||OR

Electrochemical Series

List of E noughts in order

Non Rechargeable Cells

chemicals are used up over time and EMF drops to 0v. Single use and disposed

Rechargeable Cells

the reactions are reversible (but get less efficient over time)

Fuel Cells

cells with a continuous supply of reactants that are very efficient

Bronsted Lowry Acid

proton donor

Bronsted Lowry Base

proton acceptor

Base

insoluble

Alkali

soluble base

Strong

fully dissociates in water

Weak

partially dissociates in water

Mono/Di/Tri Protic

# of acidic hydrogen in an acid

Mono/Di/Tri Basic

# of hydroxides in a base

Kw

the ionic product of water

Lewis Acid

electron pair acceptor

Lewis Base

electron pair donor

Buffer

solution that maintains an appoximately constant pH despite dilution or addition of small amounts of acid or base

Indicators

weak acids where HA and A- are different colours

Methyl Orange

red HA, yellow A-

3.2-4.4

Phenolphthalein

colourless HA, pink A-

8.2-10

Equivalence Point

the point in a titration where the number of moles of hydrogen ions equals the number of moles of hydroxide ions

Half Equivalence Point

pH=pKa

Periodicity

the repeating pattern of chemical and physical properties of the elements

Milk of Magnesia

anti acid Mg(OH)2

Slacked Lime

limewater, used for painting and soil pH maintenance Ca(OH)2

Barium Meal

BaSO4 used in X-Rays as has a different density. Insoluble so will not be toxic in blood.

Displacement Reaction

A reaction in which a more reactive element displaces a less reactive element from an aqueous solution of the latter's ions.

Transition Metal

d block metals that have a partially filled 3d sub shell and can produce stable ions

Ligand

species with a lone pair that can donate electrons coordinately

Complex

a transition metal ion surrounded by ligands that are coordinately bonded to them

Mono/Bi/Multi Dentate Ligands

1/2/many lone pairs so can form this many coordinate bonds

Chelating Ligands

form chelates which are more stable than unidentate ligands as there is a positive entropy