2. ATOMS & MOLES
• Isotopes:
• Same number of protons
• Different number of neutrons
• So mass is affected not charge
• Relative Atomic Mass:
• Average mass of an atom of an element on a scale where an atom of carbon-12 is 12.
• Mole:
• One mole is roughly 6x10
23 particles (Avogradro’s Constant, L)
• Moles x Avogadro = no molecules
• Molar Mass = Relative Molecule Mass, Mr
• Molar Mass = gmol-1
• Moles = Mass/Mr
• Moles = Vol (dm3)/24
• Concentration:
• Moles = Conc x Vol (dm3)
• ppm = % x 10,000
• e.g. 0.000009% = 0.09ppm
Particle Charge Mass
Proton +1 1
Neutron 0 1
Electron -1 1/1840
X
A
Z
Mass
Number
Proton
Number
3. EMPIRICAL/MOLECULAR FORMULA
• Empirical Formula is the smallest integer ratio of atoms in a compound
• Molecular Formula is the actual number of atoms in a molecule
• Empirical Formula Calculations:
• Mass/Mr
• Divide by smallest number
• Ratio produced
• Molecular Formula Calculations:
• Molecular Mass/Empirical Mass
• Empirical Molecule x Step 1
4. SALTS
• Solid salts consist of a lattice of positive and negative ions
• Water of crystallisation – H2O within the lattice e.g. CuSO4.5H2O
• Double Salt:
• Contains two cations or anions
• Mix two salt solutions together forming a double salt when recrystallised
• Percentage Yield is Never 100%:
• Percentage Yield = Actual/Theoretical x 100
5. ATOM ECONOMY
• Atom Economy = Molecular mass of desired product
Sum of molecular masses of all products
• Addition reactions = 100%
• Substitution have a lower economy
• Substitution ends with 2 products, only 1 is useful
• More Sustainable/Eco Friendly = High Economy
X 100
6. ENTHALPY
• Enthalpy Change = heat transferred in a reaction at constant
pressure. In kJmol-1.
• EXO = -ve, gives out energy (e.g. oxidation)
• ENDO = +ve, absorbs energy (e.g. decomposition)
• Calorimetry:
• Heat known amount of water with fuel.
• Measure temp increase.
• Q = mcΔT
• C – specific heat capacity of water 4.18
• M – mass of solution
• ΔT – temperature change
• Q – joules (to get jmol-1 ÷ by the number of moles)
• Problems with Calorimetry:
• Heat absorbed by container & lost to surroundings
• Incomplete combustion
• Evaporation of volatile fuels
• Error = (Accuracy)/Vol. Measured x 100
• e.g. +/- 0.1cm3 pipette measures 10cm3
(2 x 0.1) / 10 x 100 = 2 %
7. ERRORS
• Error = (Accuracy)/Vol x 100
• E.g. 10cm3 using a 100cm3
• 10/100 x 100 = 10%
• Heat loss so results never 100% accurate
8. DEFINITIONS
Standard Enthalpy Change of… Definition:
Reaction
Energy transferred when the molar quantities of reactants as
stated in the equation react under standard conditions
Neutralisation
Energy change when the amounts of acid & alkali in the
equation for the reaction neutralise each other under
standard conditions
Combustion
Enthalpy change when one mole of a substance burns
completely in oxygen under standard conditions
Formation
Enthalpy change when 1 mole of the compound forms from
it’s element under standard conditions with the elements and
their compounds in their standard states
Atomisation
Energy change when 1 mole of gaseous atoms is formed from
the element under standard conditions
9. HESS’S LAW
• Total enthalpy change of a reaction is always the same, no matter
which route is taken
10. BOND ENTHALPY
• Bond Breaking – Endo +ve
• Bond Making – Exo –ve
• Average bond enthalpies aren’t exact:
• Depends on environment around the bond
• Mean bond enthalpy – energy needed to break one mole of bonds in the
gas phase, averaged over many different
compounds
• Speed of reaction is dependent on bond enthalpy
• Weaker bonds = Faster reaction rate
11. MASS SPECTROMETRY
• Vaporisation:
• Sample put into vacuum – analysed as a gas
• Ionisation:
• High energy e-s knock off other e-s (+vely charged sample)
• Electron gun used
• Acceleration:
• -v charged plate pulls sample up the tube
• Deflection:
• Magnetic field introduced – lighter atoms = deflect more
• Ions all +ve so only mass varies
• Detection:
• Atoms hit charged plate – small charge is created
Uses: Drugs Testing &
Carbon dating
12. SUBSHELLS
• S – subshell can hold 2e- (1 Orbital)
• P – subshell can hold 6e- (3 Orbitals)
• D – subshell can hold 10e- (5 Orbitals)
• F – subshell can hold 14e- (7 Orbitals)
• 4s fills before 3d
• Each orbital fills with 1e- then when all are full, they pair up
• When pairing up they slightly repel so lower IE required
• Copper ends 4s13d10
• Chromium ends 4s13d5
13. IONISATION ENERGIES
• 1st I.E. of Sodium – Na(g) = Na+
(g) + e-
• 2nd I.E. of Sodium – Na+
(g) = Na2+
(g) + e-
• Ionisation Energy:
• The energy needed to remove 1 electron from each atom in 1 mole of gaseous
atoms to form 1 mole of gaseous 1+ ions.
• Factors affecting IE:
• Nuclear Charge: More protons = more +vely charged = high the IE
• Distance from Nucleus: Further away = lower IE
• Shielding: More shields = further away = lower IE
• IE decreases down group 1
• More electrons = more shells = larger atomic radius
• Electron being removed is further from nucleus
• Less energy needed to remove outermost electron
• Lower IE
IE Graph:
• Big jump = new shell
14. PERIODIC PROPERTIES
• Atomic Radius Decreases Across a Period:
• Number of protons increases – nuclei more +vely charged
• This means electrons are held more tightly
• So the atomic radius decreases
• Ionisation Energy Increases Across a Period:
• Atomic Radius – further the outer shell electrons are from nucleus = lower IE
• Nuclear Charge – more protons = higher IE
• Electron Shielding - more shells = more shielding = e- being lost further from nucleus = lower IE
• Drop in IE between Gr. 2 & 3:
• Due to extra sub-shell
• E- being lost is further from nucleus
• Drop in IE between Gr. 5 & 6:
• Due to electron repulsion
• One orbital will have 2e-, so these e- will slightly repel each other
• This means IE is lower
15. MELTING POINTS
• Periods 2 & 3 show similar trends in melting points
• Metals (Li, Be, Na, Mg & Al):
• Melting points increase across the period as the bonding strength increases
• Due to metal ions having more delocalised e- & decreased ionic radius
• This means the ions have a high charge density, this attracts the ions more strongly
• Elements with Macromolecular structure:
• Have strong covalent bonds linking all atoms together
• Carbon & silicon have highest melting points
• Simple Molecules (N2, O2, F2, P4, S8 & Cl2):
• Melting point depends on london forces
• London forces are weak and require less energy
• More atoms in a molecule = stronger VdW so S8 has higher melting point than Cl2
• Noble gases have lowest melting point because they exist as individual atoms
(monatomic)
16. IONIC BONDING
• Electrostatic attraction holds cations & anions together
• Ionic crystals are giant lattices of ions
• High melting points:
• Atoms must be held together strongly
• Often soluble in water but not in non-polar solvents:
• Particles must be charged
• Ionic compounds don’t conduct electricity as a solid:
• Can only carry a charge when dissolved or molten
• Common compounds:
• Carbonate CO3
2-
• Nitrate NO3
-
• Sulphate SO4
2-
• Ammonium NH4
+
17. IONIC BONDING
• Ions are smaller that atoms for metals but larger for non-metals:
• Metals
• Metals lose electrons when they form ions
• +ve charge of nucleus is greater than the –ve charge of the electron cloud
• When cations are formed usually the outermost shell is lost
• Outer electrons are therefore held more strongly to the nucleus
• Non-metals:
• Non-metals gain electrons when they form ions
• -ve charge of electron cloud is greater than the +ve charge of the nucleus
• Greater repulsion between electrons
• Size of ion depends on Atomic number & Charge:
• Ionic radius increases down a group – due to extra shells
• Isoelectric ions are ions with the same number of electrons, so ionic radius of isoelectronic ions
decreases as the atomic number increases – due to more protons
• Evidence for existence of ions:
• Migration of ions on wet filter paper:
• Electrolyse green copper chromate, yellow chromate ions and blue copper ions form
• Electron density maps:
• Shows that there are spaces between ions where density of electrons is zero, no shared
electrons
18. BORN-HABER CYCLES
• These cycles show enthalpy changes when a solid ionic compound is formed
from it’s elements in their standard states.
• T
• hey show 2 routes – one direct & one indirect
• From Hess’s law both routes should have the same enthalpy change
• It can show why some compounds don’t exist:
• If a lot of energy is released the compound is stable
• NaCl exists but NaCl2 doesn’t:
• Forming Na2+ requires a lot of energy
• So formation of NaCl2 is endothermic – so it requires too much
energy
• This means the formation of NaCl2 is energetically unfavourable
19. LATTICE ENERGIES & POLARISATION OF
IONS
• Theoretical lattice energies are based on ionic model:
• Experimental value – using experimental values in a born-haber cycle
• Theoretical Value – using calculations based purely on the ionic model of a lattice
• Comparing lattice energies:
• If theoretical & experimental are a close match then you can say the compounds fit the purely ionic model
• If experimental is larger than theoretical, it tells you bonding is stronger than 100% ionic
• This means that the compound must be partially covalent
• Polarisation of ionic bonds leads to covalent character in ionic lattices:
• Small cation with high charge density leads to greater polarisation
• Large anions are polarised more easily than small anions because the electrons being lost are further from the
nucleus
• For a compound to be partially covalent we need:
• High charge density small cation
• Large anion
Compound Experimental Theoretical
(assuming only ionic)
NaCl -780 -770
AgCl -905 -833
MgI2 -2327 -1944
20. COVALENT BONDING
• They are a group of molecules bonded together
• In a covalent bond, two atoms share electrons, so they have a full outer shell
• Electron density maps show electron sharing, as high electron density between the atoms
• Sigma Bonds:
• Formed when two electrons in S orbitals overlap
• Pi Bonds:
• Formed when two electrons in P orbitals overlap
• It has two parts as P orbital are dumbbell shaped, a bit above & below
• This means it’s weaker than a sigma bond & more reactive
• Dative Covalent Bonding:
• Formed when both electrons come from the same atom
21. PI & SIGMA BONDS
• Single Bond – Sigma Only
• Double Bond – Sigma & Pi
Sigma bond has a higher
electron density than a Pi
bond. Therefore Pi bond
more susceptible to
electrophilic attack
22. GIANT COVALENT STRUCTURES &
METALLIC BONDING
• Covalent bonds can form a giant molecular structure:
• Giant covalent structures have a huge amount of covalently bonded atoms
• Sometimes called a macromolecular structure
• Silicon & carbon have this structure (e.g. diamond) because they can form four strong bonds
• Silicon dioxide has a tetrahedral arrangement:
• Each silicon atom is bonded to four oxygen atoms in a tetrahedral shape forming a large crystal lattice
• Evidence for covalent bonding:
• Insoluble in polar solvents – no ions in structure
• Form hard crystals with high melting points – network of strong bonds
• Don’t conduct electricity – all electrons used to bond so no charged particles (except graphite)
• Metals have giant structures too:
• Outermost shell of electrons of a metal atom is delocalised – the electrons are free to move
• The cation is attracted to the delocalised electrons
• Forms a lattice of closely packed cations in a sea of delocalised electrons
• More delocalised electrons per atom = stronger bonding = higher melting point
• Malleable due to no bonds so layers can slide over each other
• Good thermal conductors – delocalised electrons can pass kinetic energy to each other
• Good electrical conductors – delocalised electrons can carry a current
• Insoluble – due to strength of metallic bonds
23. PROPERTIES OF MOLECULES
Bonding e.g.
Melting &
Boiling
Points
State at
RTP
Does solid
conduct
electricity
Does liquid
conduct
electricity
Soluble
in H2O
Ionic
NaCl
MgCl2
High Solid No Yes Yes
Simple
Covalent
CO2
I2
H2O
Low only have
to overcome VdW
or H Bonds
Usually
liquid or gas
but may be
solid (I2)
No No
Depends how
polarised the
molecule is
Giant
Covalent
Diamond
Graphite
SiO2
High Solid No (except
graphite)
/ No
Metallic
Fe
Mg
Al
High Solid Yes yes No
24. ORGANIC GROUPS
• General formula – an algebraic formula
• Empirical formula – simplest ratio of atoms of each element
• Molecular formula – actual number of atoms of each element in a molecule
CnH2n+1OH = Alcohols
CnH2n = Alkenes
CnH2n+2 = Alkanes
Homologous
Series
Prefix/Suffix Example
Alkanes -ane Propane: CH3CH2CH3
Branched
alkanes
alkyl- Methylpropane: CH3CH(CH3)CH3
Alkenes -ene Propene: CH3CH=CH2
Halogenoalkanes chloro-/bromo-/iodo- Chloroethane: CH3CH2Cl
Alcohols -ol Ethanol: CH3CH2OH
25. ALKANES
• Alkanes are saturated hydrocarbons
• Structural isomers have different arrangement of the same atoms
• Alkanes can have chain isomers
• Combustion:
• CH4 + 2O2 2H2O + CO2
• With Br2:
• Sunlight required, a Photochemical reaction occurs
• Free-radical Substitution:
• Initiation Homolytic fission occurs (sufficient energy in sunlight to photodissociate)
• Propagation Cl* + CH4 = HCl + CH3*
CH3* + Cl2 = ClCH3 + Cl*
• Termination Cl* + Cl* = Cl2
CH3* + Cl* = CH3Cl*
CH3* + CH3* = C2H6
All C5H12
26. PETROLEUM
• Crude oil is mostly made up of alkanes
• Fractional Distillation:
• Crude oil vapourised at 350°C
• Vapourised crude oil goes into fractionating column
• Largest hydrocarbons don’t vapourised so they just run off at the bottom
• As the crude oil rises up the column it cools
• Shortest hydrocarbons don’t condense, they just remain as gases
• Heavy fractions can be cracked into smaller molecules:
• Thermal Cracking:
• High temp (1000°C) & pressure (70atm)
• Produces lots of alkenes
• Catalytic Cracking:
• Uses a zeolite catalyst at a slight pressure and at high temp (450°C)
• Mostly makes motor fuels & aromatic hydrocarbons
27. FUELS & CLIMATE CHANGE
• Use of fuels produces harmful emissions:
• SO2 – produced mainly by power stations, scrubbers reduce these emissions
• They’re poisonous
• It dissolves in air moisture, they’re converted to sulphuric acid, causing acid rain
• CO & Hydrocarbons – caused during incomplete combustion
• CO is poisonous
• Unburned hydrocarbons are very fine & cause breathing problems
• Nitrogen Oxides – produced by vehicle engines
• Add to acid rain problem
• NO2 causes breathing problems
• In sunlight, NO2 reacts to produce ground-level ozone
• Greenhouse Gases:
• Earth absorbs radiation from sun & re-emits it as IR radiation
• GG’s in atmosphere absorb IR radiation & re-emit it towards earth again
• Climate Change:
• Melts polar ice caps
• Rising sea levels
• Famine & Droughts
28. ALKENES
• They’re unsaturated hydrocarbons
• Aromatic compounds – e.g. benzene
• Aliphatic compounds – e.g. ethane
• Alkenes more reactive than alkanes:
• C=C has high electron density
• C=C has a Pi & Sigma Bond
• H2C=CH2 + H2 C2H6
• Requires a nickel catalyst at 150°C
• Alkenes can react by electrophilic addition
• Hazard: - anything that can cause you harm
• Risk: - the chance that what you’re doing will cause you harm
29. ELECTROPHILIC ADDITION OF
ALKENES
• Test for alkenes – add Bromine Water & it should decolourise
• Product of this reaction is mainly bromoethanol
• Due to the carbocation reacting with OH- instead of Br-
• Ethene + Br2
• Initially Br2 not polar
• C=C induces polarity in Br2
• Br2 bond breaks heterolytically
• +ve intermediate is missing an e- (Carbocation)
• Prop-1-ene + HBr
• 2O more stable, because alkyl groups are slightly e- releasing – INDUCTIVE EFFECT
30. ISOMERISM
• Alkenes have stereoisomerism, due to the C=C which can’t
rotate
• We use E/Trans or Z/Cis to show the molecule’s structure
• Each carbon has to have 2 different groups attached
• Highest atomic number takes priority
Z/CisE/Trans
