Proton – +1
Electron – weighs 1/1836 -1
Neutron – changes for isotopes (neutral)
Isotopes do not change the atomic number, but increase the mass number because of the increase in neutrons,
This will change the relative atomic mass depending on spread of isotopes.
Mass spectrometry, sample in gaseous state vaporised, and bombarded with electrons, forming positive ions which are accelerated and passed by a magnet which then splits apart different strengths, forming a graph showing abundance of different species.
Can find isotopes
Relative atomic mass can be calculated
First ionisation energy
The amount of energy required to remove one mole of electrons from each atom in the gas phase to form a singly positive ion.
Second ionisation energy
The amount of energy required to remove the next electron from an atom.
Requires more energy after the first, since this will usually lead to open electrons, or electrons left in unfilled shells.
Jump to between shells requires a lot more energy, and shows large peak on graph
Across periodic table generally increases but then drops when starting a new row because, outer shell contains one electron an so is not held well due to nuclear shielding
*Slight dips occur at between elements which have half filled shells, since the next element will have an extra e- on top of the half filled shell which is out of placed, and the element is more stable without it.
S shell contains 2
P shell contains 6
D shell contains 10
*the d shell will not fill until the next s shell fills (3p 4s 3d 4p)
(ELECTRONIC STRUCTURE MODEL SHEET)
(ATOMIC ORBITAL SHAPES SHEET)
The energy change per mole of gain in electrons to form an anion in the gas phase.
First affinity generally negative (exothermic)
The second is always (positive) (endothermic)
Insoluble and covalent as usual
Cross spectator ions out (on both sides)
Moles = mass/Mr(Ar)
Moles = Molarity x volume x 10-3
Multiply/divide if solution taken from larger amount or diluted
Back-titration (unreacted method)
Calculate moles of substance used on reactants (usually alkali)
Multiply to get value of entire flask etc, e.g. if 25cm used from 250 then x 10
Work amount of moles of (acid) used altogether using Molarity x concentration
Take away amount made used altogether – amount used in reaction
This will give moles of (acid) in the equation to make up solutions, so take ratio if necessary
Calculate mass using moles x mr
Work out percentage of actual/original x 100
Bonding and structure
+ve ions are cations attracted to -ve (metals)
-ve ions are anions attracted to +ve (non-metal)
When dissolving or hydration, six water ligands by dative covalent bonding
Ionic bonding – low Ie energy
– Non-metal has high electron affinity (electron gain energy)
– Metal forms large ions with low charge
– Non-metal forms small ions of low charge
If cation is small/high charge then it is very polarising and has high charge density
If anion is big then it is polarisable
Lattice enthalpy is the measure of strength of an ionic substance, can be used for solubility.
Lattice enthalpy is the energy change per mole for the process
M+ + X- = MX
Ignores polarisation so actual covalence can be calculated
Covalence increases lattice energy
P/S orbs head on lap for sigma bonds
P orbs side lap for pi bond
A covalent bond is polar if there is a large difference in electronegativity
Electronegativity is the measure of how strongly an atom attracts electrons when in a covalent bond.
Two bonded pairs – linear 180
Two bonded two lone – bent linear 104
Three bonded pairs – trigonal planer 120
Three bonded, one lone pair – pyramidal 107
Three bonded, two lone pairs – t shaped two 90 and one 180
Four bonded pairs tetrahedral – 109.5
Five bonded pairs – trigonal bipyramid 3 are 120 and two 90
Six bonded pairs – octahedron 90
Two double bonds linear 180
One double, two single – trigonal planer 120
Two double bonds, two single – tetrahedral 109.5
Two double bonds, one lone pair – bent linear 104
Noble gases increase in temperature for mpt
Fluorine pale yellow gas
Chlorine greenish gas
Bromine brown volatile liquid
Iodine dark grey lustrous solid
The hydrogen halides are very soluble
Produce strong acids
HF has strongest bond, and decreases down the group
Have high ionisation energies
Produce ppts with Ag+ ions, of which chloride is soluble with NH3, bromide with conc. NH3 and iodide insoluble.
React with conc. sulphuric acid since halogens are reducing agents
– Chlorides will produce –hydrogen sulphates and HCL the chlorine cannot be reduced further
– bromides are stronger reducing agents since it is bigger and so loses its e easier so will produce HBR and –hydrogen sulphates, and the HBR further produces bromine and water and sulphur dioxide
– Iodide is even bigger and so further reduces conc. H2SO4, this produce the above products and then the HI further produces, I2 and H2S and water, I2 and S and water, I2 and SO2 water. I3- can also be produced in the excess of I2
All halogens are strong oxidising agents, and decreases down the group, so Cl can oxidise Br and so on.
oxo-anions – stuff with oxygen like- carbonates or hydroxides
Unit 2 – Organic and energetic
?H is the heat change within a chemical reaction under standard conditions (atmospheric pressure around 298K)
Standard enthalpy of combustion is the enthalpy change when one mole of a substance is completely burned in oxygen, under standard conditions
Standard enthalpy of formation is the enthalpy change when one mole of a compound is formed from its elements in their standard states, under standard conditions
Standard enthalpy of neutralisation is
Exothermic graph starts high and goes down visa versa for endothermic
*Graphite more thermodynamically stable than diamond so diamond not in standard state
Specific heat capacity – is the heat required to increase the temperature of 1g of substance by 1K
Heat transfer = mass x specific heat capacity x temperature change
?H = kjmol-1 (after dividing by moles, kj before dividing)
Experiment – solutions make one excess but accurate, and then divide by accurate moles for ?H
First law of thermodynamics – energy cannot be created or destroyed but only changed from on form to another
Hess’s Law – The enthalpy change for a reaction is independent of the route by which the reaction is achieved, provided that the pressure and physical states of the reactants and products are the same in each case
Enthalpy of dissociation energy – enthalpy change when mole of a gaseous substance is broken up into free gaseous atoms
Can also be called and enthalpy for covalent bonds
Carbon can catenate, form bonds with itself
homologous series – similar chemical properties, gradual variation, formula
Nucleophiles – species which seek out positive centres and must have a lone pair if electrons to donate to form a covalent bond
Electrophiles – species which seek out negative centres and must be capable of accepting a lone pair of electrons to form a covalent bond
Aliphatic – normal
aromatic – contain rings do not obey rules
Free radical (can form from breaking bonds like Cl2) homolytic fission
Calculating energy per unit – moles gives g; divide kj by moles to get kgperg then times by density for calorific value (kjcm-3)
A tertiary iodide will be the most reactive because the C-I bond is the weakest and is on the third carbon.
Aldehydes and keytones – CnH2nO (keytone cant have H as second R group)
(INSERT SYNTHETIC PATHWAY SUMMARY)
Kinetic stability means that the reactants are thermodynamically unstable but do not have enough energy to react to react and so are kinetically stable.
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