Chemistry Overlap
Amount of substance
Bonding
Chemical equilibria Kc
Energetics
PV=nRT
Empirical and Molecular formula
Mol + 6.02 x 10^23
Balanced equations and calculations RP1
Ar/Mr
Bonding and physical properties
Shapes of molecules and ions
Metallic bonding
Bond polarity
Nature of Covalent/Dative bonds
Forces between molecules
Ionic bonding
Chemical equilibria + Le Chatelier
Equilibrium constant Kc for homogenous systems
Application of Hess's law
Bond enthalpies
Calorimetry RP2
Enthalpy changes
n=moles
R=Constant (8.31 J K-1 mol-1)
V=Volume (m3)
T=Temperature (K=+273)
P=Pressure (Pa)
P=nRT/V
V=nRT/P
n=PV/RT
T=PV/nR
mol unit of measurement used to measure amount of substance. One mol contains 6.02x10^23 this number is known as the avogadro constant.
Mr = sum of Ar
Involves electrostatic attraction between oppositely charged ions in a lattice
Sulfate, hydroxide, nitrate, carbonate and ammonium
A single covalent bond contains a shared pair of electrons
Multiple bonds contain multiple pairs of electrons
A co-ordinate (dative covalent) bond contains a shared pair of electrons with both electrons supplied by one atom
Metallic bonding involves attraction between delocalised electrons and positive ions arranged in a lattice
4 types of crystal structure
metallic
macromolecular (giant covalent)
ionic
molecular
The structures of the following crystals as examples of these four types of crystal structure
graphite
ice
diamond
iodine
magnesium
sodium chloride
Bonding pairs and lone (non-bonding) pairs of electrons as charge clouds that repel each other.
Pairs of electrons in the outer shell of atoms arrange themselves as far apart as possible to minimise repulsion.
Lone pair-lone pair repulsion is greater than lone pair-bonding pair repulsion, which is greater than bond pair-bond pair repulsion.
The effect of electron pair repulsion on bond angles.
Electronegativity as the power of an atom to attract the pair of electrons in a covalent bond.
The electron distribution in a covalent bond between elements with different electronegativities will be unsymmetrical. This produces a polar covalent bond, and may cause a molecule to have a permanent dipole
Forces between molecules
permanent dipole-dipole forces
induced dipole-dipole (van der Waals, dispersion, London forces)
hydrogen bonding
The melting and boiling points of molecular substances are influenced by the strength of these intermolecular forces.
The importance of hydrogen bonding in the low density of ice and the anomalous boiling points of compounds.
Reactions can be endothermic or exothermic.
Enthalpy change (∆H) is the heat energy change measured under conditions of constant pressure
Standard enthalpy changes refer to standard conditions, i.e. 100kPa and a standard temperature (eg ∆H298Ɵ)
Calorimetry
RP2
q=mc∆T
q=heat change in a reaction measure in
m=mass of substance measured in
∆T=Change in temperature measured in Kelvin
c=specific heat capacity measured in
Measurement of an enthalpy change
Reaction enthalpy is independent of the route taken
Mean bond enthalpies not accurate
Le Chateliers principle states that if a dynamic equilibrium is disturbed by a change in its conditions the equilibrium will counteract the change by increasing the rate of the forward or backwards reaction.
Dynamic equilibria is for reversible reactions and means that concentration of both reactants and products are the same, can only happen in a closed system
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Kc is equilibrium constant
Temperature change can alter Kc
Value for Kc only valid for certain temperature
If you change the temperature of the system you will also change the equilibrium concentrations of the product and reactants so Kc will change
Energy needed to break bond depends on environment that its in
=average energy needed to break a certain type of bond over a range of compounds