This document provides information on chemistry topics including shapes of molecules, carbon structures, polar bonds, intermolecular forces, solubility, redox reactions, group 2 elements, flame tests, the halogens, indicators, kinetics, chemical equilibria, alcohols, oxidation of alcohols, haloalkanes, and nucleophilic substitution reactions. Key concepts covered include VSEPR theory, the three allotropes of carbon, electronegativity, types of intermolecular forces, factors affecting solubility, rules for oxidation numbers, reactions of group 2 elements, uses of flame tests, properties of the halogens and halides, common acid-base indicators, Maxwell-Boltzmann distribution

1. M O D U L E 2
CHEMISTRY
EDEXCEL

2. SHAPES OF MOLECULES
Valence Shell
Electron Pair
Repulsion Theory
Or VSEPR
Lone Pair: Lone Pair > Lone pair: Bond Pair > Bond Pair: Bond Pair
Electrons repel to a
point of maximum
separation to give
minimum repulsion

3. CARBON STRUCTURES
• Carbon has 3 allotropes:
• Graphite, Diamond & Fullerenes
• Diamond:
• Each carbon atom is covalently bonded with
sigma bonds to four other carbon atoms
• This creates a tetrahedral shape – crystal
lattice structure
• Properties:
• High melting point (3700K)
• Extremely hard
• Vibrations can travel through the lattice – it’s
a good thermal conductor
• Can’t conduct electricity
• Won’t dissolve in any solvent
• Graphite:
• Carbon atoms arranged in sheets of flat
hexagons
• The 4th outer electron of each carbon atom is
delocalised
• Properties:
• Weak bonds between layers – so sheets slide
over each other – used as a dry lubricant
• There are delocalised electrons – can
conduct electricity
• Less dense than diamond but strong &
lightweight
• Strong covalent bonds – high melting point
• Insoluble in all solvents
• Fullerenes:
• Can form hollow balls or tubes
• Each carbon bonds to 3 others
• There are delocalised electrons
• Can conduct electricity
• They’re nanoparticles
• Buckminsterfullerene C60 sphere shaped
• Many are soluble in organic solvents & form
bright colours
• Can be used to deliver drugs to specific cells
of the body
• Nanotube:
• Single layer of graphite rolled into a tube
• They can conduct electricity
Allotrope:
different forms of
the same element
in the same state

4. POLAR BONDS
• Electronegativity – the ability to attract the bonding electrons in a
covalent bond
e-s spend most of their time near F
• F most electronegative & Cs least electronegative
• Not all molecules with polar bonds are polar molecules
• Trichloromethane - Polar
• But, Tetrachloromethane – not Polar
• Stronger Bonds = Shorter Bonds
E.g. CsF – Large
Cation & Small
Anion
Partially
Covalent
Al3+ - Small & High Charge
So high charge density
Very polar due to F
being
electronegative.
E- in bond shared
unevenly
E.g. H - H
Electronegativity
measured on the
Pauling scale F = 4

5. INTERMOLECULAR FORCES
• VdW – London Forces
• Causes all atoms & molecules to be attracted to each other
• When e-s are on one side of an atom they form a temporary dipole
• This dipole causes another dipole in the neighbouring atom
• Larger SA/molecule/chain length = Stronger
• Dipole – Dipole
• Charges on polar molecules cause weak electrostatic attractions between
molecules
• H-Bonding
• Requires hydrogen covalently bonded to N, O, F
• NOF are very electronegative, they attract the electrons from H
• Bond is polarised & H has a high charge density
• H atoms form weak bonds with lone pairs of electrons on NOF of other
molecules
VdW < Dipole - Dipole < H-Bonding

6. SOLUBILITY
• To dissolve:
• Bonds in substance have to break
• Bonds in solvent have to break
• New bonds have to from between substance & solvent
• Non-polar solvents will dissolve non-polar solutes
• Both have VdW forces – so form similar bonds with each other
• Polar solvents (e.g. Water) will dissolve ionic substances
• Ions are attracted to oppositely charged ends of water
• Ions pulled away from lattice by H2O, called hydration
• Some ionic substance don’t dissolve because the bonding is too strong
• Not all molecules with polar bonds will dissolve, halogenoalkanes don’t
• Bond between C-X is weaker than O-H in water
• Alcohols also dissolve in polar solvents
• Polar O-H bond in an alcohol attracted to O-H in water
• H bonds form between lone pair on oxygen atoms and the hydrogen atoms
• Longer carbon chain on alcohol = less soluble
Like usually dissolves like

7. REDOX
• Oxidation – lose e-
• Oxidising Agents – gain e- & gets reduced
• Reduction – gain e-
• Reducing Agents – lose e- & gets oxidised
• Rules:
• Group 1 = +1
• Group 2 = +2
• Hydrogen = +1 (-1 in Metal Hydrides)
• Fluorine = -1
• Oxygen = -2 (-1 in Peroxides)
• Chlorine = -1 ( not in F & O2)
• Uncombined Element = 0 e.g. Cl2
• Balancing Redox Equation:
• Balance O2 by adding H2O
• Balance H2 by adding H+
• Balance charge add e-

8. GROUP 2
• With water:
• Produces a hydroxide e.g. Mg(OH)2
• Reacts more quickly down group due to low IE
• Burn in oxygen:
• Produces a oxide e.g. MgO
• React with chlorine:
• E.g. MgCl2 forms – white solid
• Oxides & Hydroxides are bases:
• Oxides react with water producing metal hydroxides
• The hydroxide ions make these solutions strongly alkaline
• MgO reacts slowly & it’s hydroxide isn’t very soluble
• Oxides form more strong alkaline solutions down the group due to the hydroxides
being more soluble
• Hydroxide more soluble down group 2
• Sulphate less soluble down group 2 (Barium Sulphate is insoluble)
• MO(s) + H2O(l)  M(OH)2(aq)
• M(OH)2(S) + H2O  M2+
(aq) + 2OH-
(aq)
• MO(s) + 2HCl(aq)  MCl2(aq) + H2O(l)
• M(OH)2(aq) + 2HCl(aq)  MCl2(aq) + 2H2O(l)
Reaction with acid
Reaction with water

9. FLAME TESTS
• Lithium – Red
• Sodium – Golden-Yellow
• Potassium – Lilac
• Calcium – Brick Red
• Strontium – Crimson Red
• Barium – Pale Green
• Copper – Blue/green
• Caesium – Blue
• Rubidium - Red
• Use Nichrome wire
• Clean wire with HCl
• Dip into compound
• Hold it over non-luminous flame
Electron Transition:
Due to electrons absorbing
energy and moving to a higher
energy level. They emit energy
(in the form of light) when they
fall back down

10. HEATING NO3’S & CO3’S
• Group 1 Nitrates:
• 2KNO3 = 2KNO2 + O2
• Group 2 Carbonates:
• CaCO3 = CaO + CO2
• Group 2 Nitrates:
• 2Ca(NO3)2 = 2CaO + 4NO2 + O2
• Group 1 Carbonates are too
thermally stable to
decompose
• Except Li2CO3
• This forms Li2O & CO2

11. THERMAL STABILITY CO3 & NO3
• More thermally stable down group 2
• E- less distorted
• Larger ionic radius
• Smaller Cation = More Polarising
E-s are pulled
towards the
cation breaking
the O-C bond

12. THE HALOGENS
Colour in Water Colour in Hexane
Chlorine Virtually Colourless Virtually Colourless
Bromine Yellow/Orange Orange/Red
Iodine Brown Pink/Violet
• Electronegativity & Reactivity Decreases down the Group
• Melting & Boiling Points Increase down the Group
• Disproportionation with Alkalis:
• Cold:
X2 + 2NaOH  NaXO + NaX + H2O
X2 + 2OH-  XO- + X- + H2O
• Hot:
3X2 + 6NaOH  NaXO3 + 5NaX + 3H2O
3X2 + 6OH-  XO3
- + 5X- + 3H2O
• Oxidise Metals:
• F2 & Cl2 are the strongest oxidising agents so we get oxidise Fe  Fe3+
• Br2 is slightly weaker oxidising agent so we get Fe2+ & Fe3+
• I2 is the weakest so only Fe2+ forms
• Oxidise Non-Metals:
• e.g. 4Cl2 + S8  4S2Cl2

13. THE HALIDES
• Reducing power increases down the group – by losing an electron
• KF/KCl with H2SO4:
Forms KHSO4(s) & HF/HCl(g)
• KBr with H2SO4:
Forms KHSO4(s) & HBr(g)
HBr then reacts with H2SO4 – Forms Br2(g), SO2(g), H2O(l)
• KI with H2SO4:
Forms KHSO4(s) & HI(g)
HI then reacts with H2SO4 – Forms I2(g), SO2(g), H2O(l)
HI then reacts with SO2 – Forms H2S(g), I2(s), H2O(l)
• Hydrogen Halides are Acidic Gases:
• Reacts with NH3(g) forming white fumes
• Blue Litmus  Red
• Halide Ions are Displaced by more Reactive Halogens:
Cl2(aq) + 2Br-
(aq)  2Cl-
(aq) + Br2(aq)
Br2(aq) + 2I-
(aq)  2Br-
(aq) + I2(aq)
• Reaction with Silver Nitrate (AgNO3):
• Fluoride – No Precipitate
• Chloride – White Precipitate - Dissolves in
Dilute NH3(aq)
• Bromide – Cream Precipitate – Dissolves in
Concentrated NH3(aq)
• Iodide – Yellow Precipitate – Insoluble in NH3(aq)
• Silver Halides React with Sunlight:
• 2AgBr  2Ag + Br2
Purple

14. INDICATORS
• Litmus paper:
• Red to Blue
Acid Alkali
• Methyl Orange:
• Yellow to Red
Alkali Acid
• Phenolphthalein:
• Colourless to Pink/Red
Acid Alkali
• Starch:
• Black to Colourless
I2 I2 + S2O3
2-
• Lead Acetate:
• H2S White to Black
• K2Cr2O7:
• SO2 Orange to Green
• Ammonia:
• Hydrogen Halide White Smoke

15. KINETICS
• Maxwell-Boltzmann Distribution:
Total area under
curve = number
of particles
Catalyst provides
alternative route
for a reaction
with a lower Ea.
10oC temp
increase = 2 x
RoR
These particles have
sufficient Ea to react
• When Temp is Increased:
• Particles have more energy
• Greater number of particles have
the Ea
• More successful collisions occur
per cm3
• Faster rate of reaction

16. CHEMICAL EQUILIBRIA
• Reversible Reactions can Reach Dynamic Equilibrium:
H2(g) + I2(g)  2HI(g)
• Concentration:
• Increase reactant = more product
• Increase product = more reactant
• Pressure:
• Increasing it shifts to side with fewer gas molecules
• Decreasing it shifts to the side with more gas molecules
• Temperature:
• Increasing temp = shifts in the endothermic direction
• Decreasing temp = shifts in the exothermic direction
2SO2(g) + O2(g)  2SO3(g)
Increase Conc of Products
Increase Conc of Reactants
Increase Pressure
-197 kJmol-1
Decrease Temperature
Increase Temperature

17. ALCOHOLS
• 1o – Functional group attached to a carbon which has 1 carbon attached to it – Least Reactive
• 2o - Functional group attached to a carbon which has 2 carbons attached to it
• 3o - Functional group attached to a carbon which has 3 carbons attached to it – Most reactive
• Producing Halogenoalkanes using Phosphorus Halides:
• 3ROH + PX3  3RX + H3PO3
• Properties:
• Liquid at room temp – due to H bonds between molecules
• Low Volatility – due to H bonding
• Soluble in H2O – due to H bonding – but less soluble as chain length increases
• Alcohols & Na  Alkoxides:
• Longer chain length = less reactive
• 2CH4O + 2Na  2CH3CH2O- Na+ + H2
Methanol Sodium Methoxide
• Reaction with PCl5 – test for OH group:
• Misty white fumes given off when NH3(g) added
• C2H5OH(aq) + PCl5(s)  C2H5Cl(aq) + HCl(g) + POCl3(aq)
• HCl(g) + NH3(g)  NH4Cl
PI3 & PBr3 are made in situ using Red
Phosphorus

18. OXIDATION OF AN ALCOHOL
• Partial oxidation – limited Na2Cr2O7 & dilute H2SO4
• Reflux – Conc H2SO4 & excess Na2Cr2O7
• Colour changes from orange to green
• Reflux:
• Allows reactions to happen at highest temperature without
loss of product or reactants
• Water goes in at bottom to ensure constant cooling
• Separation of Products:
• Distillation
• Alcohol into Halogenoalkanes:
• Add PCl5
• Reflux + NaBr +50% H2SO4
• I2 & moist red phosphorus  Iodoalkanes, react
with alcohol
LiAlH4 to go from Acid
to 1o or Aldehyde

19. PREPARATION OF ETHANAL BY
OXIDATION OF ETHANOL
• 50cm3 H2O in 500cm3 flask, add 17cm3 of conc H2SO4 & anti-bumping
granules
• Put flask in distillation apparatus. Still head has tap funnel & receiving
flask in ice-bath (ethanal low B point – avoid evap)
• Dissolve 50g of Na2Cr2O7 in 50cm3 H2O in small beaker. Add 40cm3 of
ethanol. Stir thoroughly
• Heat flask until boils, remove heat. Run alcohol/dichromate solution
slowly into flask, mixture becomes green, takes 20 mins. Maintain
gentle boiling
• Aqueous solution of ethanal collects in receiver.

20. HALOGENOALKANES
• Reaction with KOH(aq):
• Heat under reflux with KOH(aq) giving an alcohol
• Reaction with KOH in ethanolic solution:
• Eliminates hydrogen halide forming alkene
• Reaction with Conc. Ammonia in Ethanol:
• Heat & Pressure with Conc NH3 to produce amines – Nucleophilic Addition
• Preparation of 1-Bromobutane:
• 30cm3 of H2O, 35g NaBr, 25cm3 of Butan-1-ol
• Add 25cm3 of conc H2SO4 drop by drop occasionally cool
• Reflux for 45 mins
• Distil off crude 1-bromobutane (about 30cm3)
• Shake distillate with water in separating funnel and run off lower layer of 1-bromobutane
• Add 1-bromobutane back into funnel & add half it’s vol of HCl
• Shake with Na2CO3 releasing pressure
• Run off lower layer and add granular anhydrous CaCl2, swirl until clear
• Filter into clean, dry flask and distil it. Collect fractions between 99-102oC
Uses:
• Solvents
• Refrigerants
• Pesticides
• Fire Extinguishers

21. NUCLEOPHILIC SUBSTITUTION
• SN2 – All 1o halogenoalkanes react this way
• SN1 – All 3o halogenoalkanes react this way

22. TYPES OF REACTION
• Addition – Joining 2+ molecules together forming a larger molecule
• Polymerisation – joining monomers together forming a polymer
• Elimination – small group of atoms breaks away from a larger molecule
• Substitution – one species is replaced by another
• Hydrolysis – splitting of a molecule by adding H+ & OH- from H2O
• Oxidation – reaction in which an atom loses electrons
• Reduction – reaction in which an atom gains electrons
• Redox - reaction in which electrons are transferred between 2 species
• Homolytic Fission:
• X-Y  Xo + Yo
• Forms Free-Radicals
• Heterolytic Fission:
• X-Y  X+ + Y-
• Forms cation & anion
• Electrophiles are electron pair acceptors
• They are +ve
• Nucleophiles are electron pair donors
• They are –ve

23. OZONE – O3
• O2 + hv  O + O
• O + O2  O3
• O2 + O  O3
• CCl3F  CCl2F + Cl
• Cl + O3  O2 + ClO
• ClO + O3  2O2 + Cl
• 2O3  3O2 Cl is the catalyst
UV Radiation
Nitric Oxide breaks down Ozone too
NO

24. INSTRUMENTAL ANALYSIS
• Vapourisation – Ionisation – Acceleration – Deflection – Detection
• Organic Molecule Detection:
• E.g. C2H5OH can be vapourised & ionised – C2H5OH+ - appears at 46m/e (parent ion
peak)
• When fragmented charges allow us to be what group is present
• IR Spectroscopy:
• All bonds stretch and bend naturally
• Polar bonds change polarity and absorb that frequency of IR as they vibrate
• O – H 3600 wavenumber/cm-1 smooth u curve
• C = O 1740 wavenumber/cm-1 steep v curve Aldehydes
• O – H 3500 wavenumber/cm-1 smoother than O – H
• C = O 1690 wavenumber/cm-1 steep v curve Ketones
• C = O 1710 wavenumber/cm-1 Carboxylic Acid