M O D U L E 2

CHEMISTRYEDEXCEL

SHAPES OF MOLECULESValence Shell Electron Pair Repulsion TheoryOr 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

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

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 ChargeSo 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

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

moleculesVdW < Dipole - Dipole < H-Bonding

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

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-

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

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

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

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

THE HALOGENS

Colour in WaterColour 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

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

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 + S2O32-

• Lead Acetate:• H2S White to Black

• K2Cr2O7:• SO2 Orange to Green

• Ammonia:• Hydrogen Halide White Smoke

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

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

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

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

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.

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

NUCLEOPHILIC SUBSTITUTION

• SN2 – All 1o halogenoalkanes react this way

• SN1 – All 3o halogenoalkanes react this way

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

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 tooNO

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