Edexcel Chemistry Unit 2 Revision Notes

Chemistry Unit 2: Notes

2.1-Shapes of molecules and ions

Electron-pair repulsion theory: The shape of a molecule or an ion depends on the number of electron pairs that surround the atom. The electron pairs repel each other so stay far apart from each other.

Linear: 2 electron pairs, 180˚ bond angles (e.g. BeCl2, CO2) Trigonal Planar: 3 electron pairs, 120˚ bond angles (e.g. BCl3) Bent: 2 electron pairs, 1 lone pair,

Tetrahedral: 4 electron pairs, 109.5˚ bond angles (e.g. CH4, NH4+ )

Trigonal Pyramidal: 3 electron pairs, 1 lone pair, 107˚ bond angles between electron pairs (e.g. NH3)

Bent: 2 electron pairs, 2 lone pairs, 104.5˚ bond angles between electron pairs (e.g. H20)

Trigonal bipyramidal: 5 electron pairs, 90˚ and 120˚ bond angles (e.g. PCl5) Octahedral: 6 electron pairs, 90˚ and 180˚ bond angles (e.g. SF6)

Lone pairs and double/triple bonds repel more than electron pairs hence the differences in angles even with 4/ 3 bonds.

Alkanes: Tetrahedral due to all bond angles being 109.5˚ due to carbon forming 4 bonds.

Alkenes: Bonds around the double bond are trigonal planar but due to the double bond the bond angles can range from 121˚-118˚.

Alcohols: Hydrocarbon chain is tetrahedral so bond angles are 109.5˚. C-O-H bond is a bent molecule so is 104.5˚.

Carboxylic acids: -COOH group around the carbon is trigonal planar so 120˚. Haloalkanes: Tetrahedral as halogen bonds don’t affect bond angles. Same as alkanes

Carbonyls: CHO/C=0 bonds around the carbon are trigonal planar (120˚). Diamond: Tetrahedral due to 4 carbon electron pairs being formed.

Graphite: Trigonal planar as only 3 carbon electron pairs are formed. Good conductor due to weak London forces

Fullerene: Spherical molecule made up of about 60 carbon atoms which can dissolve in petrol and is a good conductor.

Nanotubes: Tube-like cage structure made up of 12 membered rings that can be used to carry drugs to target body cells.

2.2-intermediate bonding and bond polarity

Electronegativity: A measure of the attraction of an atom in a molecule for a pair of electrons in a covalent bond.

Fluorine is the most electronegative as it has the least shielding whilst still having a large number of protons so has a greater pull.

Covalent bond: Shared electrons. Ionic bond: Loss of electrons by metal and gain of electrons by non-metal. Both are two extremes of the spectrum.

Differences between electronegativities of atoms can determine what percentage of each character (covalent or Ionic) they are. Looking at the electron cloud around the bond can also determine their character. Polar covalent bonds will have an even spread of charge whereas Ionic bonds will show a distortion in the electron cloud around one atom.

Non-symmetrical molecules tend to be polar so will produce a clear dipole (e.g. Trichloromethane) with a known polarity measured in dipole moment (Debye, D).

Symmetrical molecules are non-polar due to the dipoles cancelling. (e.g. CO2)

2.3-Intermolecular forces

Permanent dipoles: Polar molecules only. Negative dipole attracted to the Positive dipole and vice versa. 100x weaker than covalent bonding.

London forces: Electrons distributed around the nucleus change so more electrons are distributed around one end. This induces another atom that’s near it to do the same. So a weak temporary bonding forms between the two atoms. Is a weak force but strength does depend on the size of the electron cloud. More electrons=more delocalisation=more London forces.

Hydrogen bonding: Hydrogen attached to a very electronegative element (e.g. fluorine, oxygen and nitrogen.). This is the strongest intermolecular force.

Boiling temperature trends (determined by intermolecular forces)

Alkanes with increasing chain length: Increasing boiling temperature due to increasing London forces with the increase in the number of electrons.

Branching in the carbon chain: The boiling temperature decreases. This is because the side chains interfere with the packing of the molecule. Molecules can’t form many intermolecular forces.

Alcohols: Have hydrogen bonds present so have high boiling temperatures. Water has a higher boiling temperature than alcohols as it can form double the amount of Hydrogen bonds.

Hydrogen halides down the period: Fluorine is highly electronegative so has a very large boiling temperature compared to the rest which from chlorine starts low around -80˚ and then gradually increases due to increased dipole-dipole interactions from increased number of electrons.

Solubility trends

A solution is made up of a solute and a solvent. Table below shows the solubility of different substances. Ketone is unusual in that it can dissolve in both water and organic solvents as it has both of the functional groups.

Formula Description E.g.

Polar AB Linear CO

HA Single H HCl

AxOH OH at one end C2H5OH NxAy N at one end NH3 Non-polar A2 All elements O2 CxAy Carbon compounds CO2

2.4-Redox

Oxidation no.: amount of charge due to no. of electrons. Rules

Atom Oxidation no.

Elements 0

Uncombined ion It’s charge

Molecule total 0

Fluorine -1

Hydrogen +1 (except metal hydrides=-1)

Oxygen -2 (except peroxides=-1 and with

fluorine=positive no.)

Chlorine -1 (except with oxygen or

fluorine=positive no.)

Group 1, 2, 3 +1, +2, +3 respectively

OILRIG: Oxidation is loss, Reduction is gain.

Reducing agent: Reduces another substance whilst being oxidised itself Oxidising agent: Oxidises a substance whilst being reduced itself. Reduction half equation: Cl2(g) + 2e- 2Cl-(aq)

Oxidation half equation: 2I-(aq) I2(s) + 2e

-Key half equations: 02(g) +4e- 202-(s)

2H20(l) +2e- 2OH-(aq) +H2(g) 2H+(aq) +2e- H2(g)

Displacement: One substance replaces another in a reaction.

Highly Polar solids Polar organic substances Non-polar solids Ionic compound (e.g. NaCl) Non-polar liquids

Water soluble soluble insoluble Soluble (Hydration enthalpy) Hexane insoluble insoluble soluble

Polar liquids Immiscible (forms two separate layers) Non-polar liquids miscible

Disproportionation: Where a substance is both oxidised and reduced in a reaction.

2.5-Periodic Table

Trend in 1st Ionisation energy: Decreases down the group due to increasing atomic radius where the outer electrons are further away so less energy is needed to

remove an electron. Forms 2+ ions easily

Reactions- With Oxygen: Burns brightly to produce a metal oxide with increasing reactivity down the group.

With Chlorine: Solid metal chloride formed. Also has an increase in reactivity. With Water: Increasing reactivity down group. Beryllium has no reaction,

Magnesium only reacts with steam to produce a magnesium oxide and the rest react with cold water to produce metal hydroxides.

Oxides with Water: Increasing reactivity. Beryllium and magnesium only react slightly, Calcium fizzes to produce calcium hydroxide (slaking lime), and strontium and Barium react in a similar way.

Oxides/Hydroxides with Dilute Acids: Forms a salt and water. E.g. CaCl2+ H2O or Ca(NO3)2+H2O

Flame Testing: Place nichrome wire in

concentrated HCl then heat in a bunsen flame to clean. Colours of flames are shown in table. Solubility

Of hydroxides: Increases down group Of Sulphates: Decreases down group Thermal Stability

Of Nitrates: Decreases down group 1 due to weaker charge of attraction. Group 1 are more stable than group 2 apart from lithium due to larger charge.

2NaNO3 2NaNO2+O2, 4liNO3 2Li2O+4NO2+O2, 2Mg(NO3)2 2MgO+4NO2+O2 Of Carbonates: More stable as you go down the group as cations get bigger so there is a more polarising effect on the carbonate.

Common Oxidising Agents Common Reducing agents O2+4e- 2O2- M Mn++e- (metal) Cl2+2e- 2Cl- Fe2+ Fe3++e -Br2+2e- 2Br- 2I- I2+2e -I2+2e- 2I- 2S2O32- S4O62-+2e -Fe3++e- Fe2+ C2O42- 2CO2+2e -2H++2e- H2 H2O2 O2+2H++2e -Mn04-+8e- Mn2++4H2O S032-+H2O SO42-+2H++2e -Cr2O72-+14H++6e- 2Cr3++7H2O 2H2SO4+2e- SO42-+2H2O+SO2

Cation Flame Colour Lithium Red Sodium Yellow Potassium Lilac Magnesium No colour Calcium Orange/Red Strontium Red

Group7-Halogens Solubility

In Water: Chlorine forms chlorine water which is colourless, Bromine water is Orange; Iodine isn’t soluble in water as it needs I

ions so instead, it dissolves in Potassium iodide which is Brown.

In organic solvents: As halogens are non-polar, they dissolve more easily in hydrocarbon solvents than water. Chlorine= Pale yellow, Bromine=Brown, Iodine=Purple

Oxidisation reactions

Halogens are strong oxidising agents,

With metals: Form salts. 2Fe+3Cl2 2FeCl3, Fe+I2 FeI2

With Hydrogen: Decreasing reactivity down group. Chlorine explodes whereas Bromine and Iodine need 300˚C and Platinum catalyst.

With Phosphorus: Chorine forms a chloride (III) then a chloride (V).

With Iron Chloride: Green iron (II) chloride becomes oxidised to brown iron (III) chloride.

Chlorine is a stronger oxidising agent than iodine: 2KI-+Cl2 2KCl+I2

Sulphuric acid oxidises halides and halogens with increasing vigour down the group.

Hydrogen Halides are covalent but become increasingly ionic as you go up the group

Hydrogen Halides and Ammonia: NH3+HCl NH4Cl Disproportionation reactions

With cold dilute Alkali: Cl2+2OH- Cl-+ClO-+H2O With warm Alkali: 3ClO- 2Cl-+ClO3

-Testing for halogens

Silver halides: AgCl-White precipitate which is soluble in ammonia, AgBr- Cream precipitate which is soluble in ammonia, AgI- Yellow precipitate which is insoluble in ammonia.

Concentrated sulphuric acid and glass rod with ammonia: Chlorine=white fumes, white smoke. Bromine=white and orange fumes, white smoke. Iodine=White and purple fumes, white smoke.

2.6-Kinetics

Rate of reaction: Speed with which reactants disappear and products are formed for a particular reaction.

Decreased rate=Decreased yield Factors in rate of reaction

Concentration: Affects number of collisions due to change in number of particles in a given volume.

Temperature: Affects activation energy due to faster moving particles with more successful collisions.

Pressure: Affects number of collisions due to change in number of particles in a given volume.

Surface area: Affects the number of particles open to contact.

Catalysts: Changes the rate of reaction without being used up or undergoing any permanent changes by lowering the activation energy through forming an activated complex. Industrial processes rely on catalysts to save money, energy and

resources.

Collision theory: In order for a reaction to happen, colliding particles need to be in the right orientation and have enough energy. Rate of reaction depends on this. Maxwell-Boltzmann model- models the distribution of molecular energies so changes in concentration, temperature or pressure can be calculated to predict the rate of reaction. E.g. As the temperature rises, the graph flattens so there are a greater proportion of particles moving fast enough to overcome activation energy. Activation energy: Minimum energy required for a reaction to take place.

2.7-Chemical Equilibria

Dynamic Equilibria: Two opposing processes that occur at the same rate so have constant macroscopic properties.

Le Chateliers principle: Whenever a system in dynamic equilibrium is disturbed, it tends to respond in opposition to the disturbance in order to restore equilibrium. Increase in Temperature: Moves in the endothermic direction.

Increase in Pressure: Moves to the side with the fewest number of moles. Increase in Reactants: Moves to the other side as more products are formed. Increase rate.

Equilibria Reactions: N2O4 2NO2, ICl+ICl2 ICl3

Haber Process: Production of ammonia (Exothermic reaction). Atom economy= 100% as all reactants are used (Recycled).

2.8-Organic Chemistry

Functional group: -OH (Methanol, Ethanol, Propan-1-ol)

Primary Alcohol: 1 carbon attached to the carbon with the functional group on it. Secondary Alcohol: 2 carbons attached.

Tertiary alcohol: 3 carbons attached.

Combustion: Produces carbon dioxide and water

Reaction with sodium: 2Na+2C2H5OH 2C2H5O-Na++H2, effervescence, forms a

Reaction with PCl5: Reaction produces a chlorine haloalkane, POCl3 and HCl gas

(misty white fumes that turn damp blue litmus paper red).

Oxidation using acidified potassium dichromate (toxic, carcinogen): Primary alcohols when oxidised, produce aldehydes, then carboxylic acids if refluxed. Secondary alcohols produce ketones which can’t be oxidised further. Tertiary alcohols show no reaction.

When making ethanoic acid, add ethanol carefully to dichromate as it could

evaporate if added quickly. Fractional distillation is used to separate ethanoic acid from the waste product, water.

Distillation Apparatus:

Reflux Apparatus:

Halogenoalkanes

Functional group: R-Cl, R-Br, R-I

Primary Haloalkanes: 1 carbon attached to the carbon with the functional group. Most reactive as nucleophiles are more attracted.

Secondary Haloalkanes: 2 Carbons attached. Tertiary Haloalkanes: 3 Carbons attached.

Structural isomers are very common in haloalkanes. A change in halogen atom position makes a huge difference to the properties of the molecule.

Halogens react faster in substitution reactions if the halogen is attached to a branched chain.

Metal halides and concentrated sulphuric acid should not be used in the creation of a haloalkane as they are very reactive.

Reaction with aqueous KOH: Substitution reaction which creates an alcohol as :OH- is attracted to the carbocation after breaking H-Cl.

Reaction with alcoholic KOH: Elimination reaction which produces an alkene. KOH attracts H+ by acting as a base.

Reaction with silver nitrate dissolved in water: Ag+ ion reacts with the halogens to create insoluble products. (See page 6-Silver halides)

Reaction with alcoholic ammonia: ammonia acts as a nucleophile which replaces the halogen atom in a halogenoalkane.

Uses of halogenoalkanes

Anaesthetic: To put patient in a deep sleep. Chlorine= deep sleep, Fluorine-Carbon bonds= stable molecule so less toxic.

Fire retardants: Used to be in fire extinguishers but is now in circuit boards, fabrics and clothing.

Refrigerants: Liquid that changes phase to keep things cool. Also used in: Aerosol containers and Insecticides

2.9-Mechanisms

Addition: Two or more substances react to form a single product.

Elimination: Small molecule removed from a larger molecule leaving a double bond.

Condensation: Addition followed by elimination of H2O or HCl.

Substitution: One atom or group of atoms replaced by another atom or group of

atoms.

Oxidation: Process where electrons are lost. Reduction: Process where electrons are gained.

Hydrolysis: Splitting of a molecule by reaction with H2O.

Polymersiation: Joining of small molecules together in a long chain.

Homolytic fission: When a bond is broken and both atoms take a single electron

each and form two free radicals.

Heterolytic fission: When a bond is broken and both electrons only go to one atom

creating a negative ion and a positive ion. (Carbanion and carbocation.)

Free radicals: Has an unpaired electron so is extremely reactive and short lived.

E.g. ∙Cl, ∙CH3

Electrophile: Atom attracted to an electron-rich centre. Has a positive charge. E.g.

H+, Br+

Nucleophile: atom attracted to an electron-deficient centre. Has a negative charge.

E.g. :NH3-, :OH

-Classifying reagents gives clues to how a mechanism will take place. Bond polarity determines whether the centre will be rich or electron-deficient.

Nucleophillic substitution reaction: Halogenoalkanes Free-radical substitution reaction: Alkanes with a halogen.

Ozone layer: Formed by O·+O2 O3. This layer absorbs UV light from the sun but when cloud cover and chemicals build up, the spring UV light hitting the clouds causes any CFCs in them to break up and create free radicals which break down the ozone layer into O2 molecules.

2.10-Mass spectra and IR absorption

Highest m/z Value= Parent ion. Most Abundant peak= Base peak.

Peaks are produced by the fragmentation of organic molecules. IR absorption

C-H Stretching=2820-3010

O-H Stretching=Broad peak at 2500-3750 N-H Stretching=3300-3500

C=O Stretching=1680-1740 C-X Stretching=500-1400

There is a change in functional groups when alcohol is oxidised to a carbonyl or carboxylic acid.

Only molecules that change polarity when they vibrate can absorb IR radiation e.g. Greenhouse gases.

Diatomic molecules can’t absorb IR.

2.11-Green Chemistry

Bio-oil is made for fuel from pyrolysis (heating wheat without burning). Ethanol is

produced from organic waste by converting sugars using bacteria or yeast. Starch has lots of uses including adhesives, paper bonding, textile fibres, shopping bags, absorption paper, drug and pesticide encapsulation and corrugated card.

Catalysts enable a reaction to go under lower temperatures and pressures to save

energy. Catalysts need to be cheap, very active and produce no by-products.

Microwave ovens can heat reactants more economically. Continuous pulses of

radiation heats up reactants. Electric field occurs which rotates polar molecules and lines them up.

Recycling glass and aluminium saves energy from having to make more. Water as

a waste product from a reaction in a factory must be treated before being disposed to avoid contaminating the environment. Waste acidic gases are removed from the rest of the air before it goes into atmosphere.

CO2 is the most worrying greenhouse gas as it absorbs lots of IR and is abundant in the atmosphere. The other greenhouse gases aren’t as abundant.

Athropogenic factors: due to activities of human beings. Natural factors: due to natural processes on earth.

Carbon footprint: a measure of the impact on environment from how much

greenhouse gas is produced. (Measured in CO2)

Petrol isn’t carbon neutral due to the slow process of formation of crude oil.

Bio-ethanol isn’t carbon neutral either as production of biofuel requires energy.

Hydrogen isn’t carbon neutral as it creates CO2 when it is being formed from

methanol.

Cl˙+O3 ClO˙+O2

ClO˙+O3 Cl˙+2O2

So in total, 2O3 3O2 due to CFCs so more UV radiation therefore hits earth.

Titration calculations

To find percentage of a metal in an impure substance

1. Find the no. of moles using the known concentration and volume of a substance.

2. Use the molar ratio to find the no. of moles in the other reacting substance with the metal.

3. If divided by 10 fold in question, x no. of moles by 10 to get original no. of moles.

4. Original no. of moles x Mr of metal= Mass of metal in grams 5. Mass of metal/Total mass x100= Percentage of metal