Key Definitions Flashcards
ISOTOPES
Atoms of the same element with different numbers of neutrons and different masses
ATOMIC NUMBER
Number of protons and electrons in an atom
MASS NUMBER
Total number of protons and neutrons in the nucleus of an atom
IONS
Particles of an element with different numbers of electrons and protons that also have a charge
RELATIVE ATOMIC MASS
The average mass of an atom compared to 1/12 of the carbon 12 atom
RELATIVE ISOTOPIC MASS
The mass of an isotope relative to the mass of 1/12 of the carbon 12 atom
MOLE
A mole is the amount of a substance that contains as many particles as there are atoms in 12g of carbon
MOLAR MASS
The mass of 1 mole is equal to the relative molecular mass and is written in g/mol
AVOGADRO’S CONSTANT
The total number of atoms in 1 mole. There are 6.02x10^23
MOLECULAR FORMULA
The actual number of atoms of each element in a compound
EMPIRICAL FORMULA
The simplest whole number ratio of atoms of each element in a compound
RELATIVE MOLECULAR/FORMULA MASS
The total atomic masses of each element added together. Relative molecular is used for simple molecules. Relative formula is used for giant structures like ionic
STOICHIOMETRY
The ratio of an amount in moles in an equation
PERCENTAGE YIELD
The amount of substance that is made compared to the theoretical amount of substance.
LIMITING REAGENT
If a reactant is not in excess, it will be used up first, therefore it is the limiting reagent, this is normally the one with bigger amount of moles.
ATOM ECONOMY
How well an atom is used up, a perfect atom economy is 100%
STRONG ACID
When dissolved in water it releases all of its H+ protons and fully dissociates
WEAK ACID
Only releases a small amount of H+ ions when dissolved in water and partially dissociates.
BASE
Metal oxide, metal hydroxide or ammonia
Proton acceptor
ALKALI
Base that releases OH- ions when dissolved in water
SALT
Product of a neutralisation reaction where its H+ protons are replaced by a metal or ammonium ion.
ORBITAL
a region of space where it is most likely to find an electron. It is in an electron cloud that can only hold 2 electrons at a time
IONIC BOND
the electrostatic attraction between a postive and negative ion to form an ionic compound
COVALENT BOND
the strong electrostatic attraction between a shared pair of electrons and the nucei of the bonded atoms
DATIVE COVALENT BOND
covalent bond where the electrons are shared from one of the atoms that usually existed as a lone pair.
ELECTRONEGATIVITY
The ability for an atom in a covalent bond to attract the shared pair of electrons towards itself.
POLAR
In order for a molelcule to be polar, there must be differences in electronegativity between the two bonded atoms. One of the atoms will attract the pair of electrons towards itself and introduce a dipole. A polar molecule will have positive and negative regions in the molecule which is a permanent dipole
LONDON FORCES
These are induced dipole-dipople interactions that occur in all compounds when the electron moves around, it will create an instantaneous dipole which will introduce a dipole in the neighbouring molecules.
PERMANENT DIPOLE
These are the dipoles of a polar bond that will be there permanently because of the difference in electronegativity
HYDROGEN BONDS
This is the electrostatic attraction between the lone pair of electrons on a partially negative atom on one molecule and an electron deficient partially positive hydrogen on another molecule.
FIRST IONISATION ENERGY
This is the energy taken to remove one mole of electrons from one mole of gaseous atoms to make one mole of gaseous 1+ ions
SECOND IONISATION ENERGY
This is the amount of energy needed to remove an electron from one mole of gaseous 1+ ions of an element to make one mole of 2+ ions
METALLIC BONDING
The electrostatic attraction between a metal cation and a negative delocalised electron that make up the metals structure.
GIANT COVALENT LATTICE
This is a giant structure that has covalent bonds on all sides bonding them together. It’s usually carbon that makes up the lattice and can exist in many forms such as diamond, graphite or graphene.
ENTHALPY CHANGE
How the heat energy has changed and you normally do products - reaction enthalpy to work it out
LAW OF CONSERVATION OF ENERGY
It states energy cannot be created or removed, it can only be transferred between systems
EXOTHERMIC
When more energy is transferred from the system to the surroundings, so the enthalpy is negative. This is when you form bonds
ENDOTHERMIC
This is when energy is transferred from the surroundings to the system so the enthalpy change is positive and this happens when bonds are formed.
ACTIVATION ENERGY
This is the minimum amount of energy needed for a reaction to take place which breaks the bonds of the reactants so they can start reacting
ENTHALPY CHANGE OF REACTION
This is the enthalpy change when reactants as seen in the reaction equation in their standard states react under standard conditions to form products in their standard states
ENTHALPY CHANGE OF FORMATION
The enthalpy change when 1 mol of product is formed in its standard states when it’s elements react in their standard states under standard conditions
ENTHALPY CHANGE OF COMBUSTION
The enthalpy change when 1 mol of a substance reacts completely with oxygen in their standard states under standard conditions to form products in their standard states.
ENTHALPY CHANGE OF NEUTRALISATION
Enthalpy change when one mole of water is formed after an acid and a base react together in standard conditions, all reactants and products are in standard states.
SPECIFIC HEAT CAPACITY
The amount of energy required to raise 1g of a substance by 1K
AVERAGE BOND ENTHALPY
The amount of energy required to break 1 mol of a specific bond in a gaseous molecule.
HESS’ LAW
States that a reaction can take place in two alternate reaction pathways as long and as long as the start and finishing conditions are the same, the average bond enthalpies will be the same.
COLLISION THEORY
It states that in order for a successfull reaction to take place, two reacting particles must collide successfully at the correct orientation
HOMOGENOUS CATALYST
This is a catalyst that has the same states as the reactants and will react to form an intermediate. E.g. esters use H2SO4 as a catalyst
HETEROGENOUS CATALYST
This is a catalyst that has different physical states than the reactants in the reaction. They will absorb the reactants, catalyse them then desorb them.
DYNAMIC EQUILIBRIUM
To be in dynamic equilibrium, the forward reaction has to be the same rate as the reverse reaction and the concentration of reactants and products don’t change.
LE CHATELIERS PRINCIPLE
He stated that when a system in equilibrium undergoes a change, the equilibrium will shift and readjust to minimise that change.
