Edexcel as Chemistry Unit 2

Edexcel AS Chemistry–Unit 2 Bonding and intermolecular forces 

Electron pair repulsion theory

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Bond angles

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Linear 180°

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Triganol Planar 120°

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Tetrahedral 109.5°

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Triganol Bipyramidal 90° and 120°

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Octahedral 90°

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Lone pairs of electrons count toward an area of electron density but as they repel to a much greater extent they reduce the bond angle by 2.5°

Bond length – the stronger the bond the shorter the length

Electronegativity

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The ability to attract the electrons in a covalent bond

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Ionic property of covalent bonds

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Same applies the other way around

Differences in electronegativity cause permanent dipoles

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Polar substances dissolve in water as they disrupt hydrogen bonding lattice

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Non-polar substances will not dissolve in water

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Polar bonds break first in organic chemistry

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Polar bonds are always heterolytic fission

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This is because there is not enough energy to make the non-polar substances break the bonds within the polar water molecules

Produces one electrophile and one nucleophile

London forces

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Exist between all atoms

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Weak force relative to size of electron cloud

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Random fluctuation will induce instantaneous dipoles

Hydrogen bonding

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Formed when hydrogen bonds to a highly electronegative atom with an unpaired electron

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Formed at 180°

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Strongest intermolecular bonds

Physical properties from bonding

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Allotropes of carbon

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Diamond – 4 bonds, giant covalent structure, requires high energy to break as there are numerous strong bonds. No conduction of electricity

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Graphite – 3 bonds in flat sheets, makes it very slippery and useful as a lubricant. Weak bonds between flat layers are london forces. Delocalised electron so conducts electricity

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Fullerenes – Ball shaped (Bucky ball) so strong due to distribution of bonds, conducts also

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Nanotubes – Hollow tubes of carbon atoms, very strong due to rigid shape and conducts electricity

Organic chemistry

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Longer chains have more electrons so greater london forces

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More branching reduces boiling points as chains can’t get as close to each other so london forces are weaker

Inorganic Chemistry 

Ionisation energies decrease down group 2

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This is because the atomic radii increase and the shielding stays the same so the effective nuclear charge is the same but further away

Group 2 reactions:

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Group 2 + Oxygen –> Metal Oxide (solid)

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Group 2 + Water –> Hydroxide + Hydrogen gas

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Group 2 + Chlorine –> Metal Chloride (solid)

Group 2 Oxide reactions

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Oxide + Water –> Hydroxide

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Oxide + Acid –> Chloride/Nitrate + H2O

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Group 2 Hydroxides will also do this, 2H2O results

Group 2 sulphates

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Solubility decreases down group

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Test for sulphates: add to Barium Chloride, will give white precipitate

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Group 2 hydroxides become less soluble down the group

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Thermal stability of group 1 and 2 carbonates

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Group 1 – stable, they don’t decompose from heating

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Group 2 carbonates + heat –> Group 2 oxide + CO2

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Stability of nitrates and carbonates increase down the group

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This is because oxides are more stable, and the first elements are more polarising so they attract the oxygen away from the carbonate

Group 1 and 2 nitrates:

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Group 1 decompose to nitrite and oxygen

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Flame tests

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Caused by electron being excited to a higher energy level, when it decreases energy level the extra energy is released as a photon

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Fluorine is the strongest oxidising agent as it has the strongest attraction from the nuclear charge where it is so small

Disproportionation with alkalis

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Cold will lead to something like 2NaOH + Cl2 –> NaCl + NaClO + H2O

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Hot forms the halate: 2NaOH + Cl2 –> NaCl + NaClO3 + H2O

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So for bromine: cold leads to bromide, but if it’s hot you get a bromate

Iodine thiosulphate titrations

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An oxidising agent will react with iodine ions

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These iodine ions will then react with thiosulphate

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Colour goes from yellow to colourless with just thiosulphate

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(Physics!)

Halogens are oxidising agents; their strength decreases down the group

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Group 2 decompose to oxygen, nitrogen and oxide

Adding starch makes it go black as there are still iodine ions, as they react with thiosulphate it goes colourless

Test for halides

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Add silver nitrate and colour of precipitate

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White = Cl

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Cream = Br

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Yellow = I

Confirmation test using ammonia

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Cl will dissolve with dilute ammonia

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Br will dissolve with concentrated ammonia

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I will not dissolve

Reaction rates 

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Increased by increasing:

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Concentration

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Pressure

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Surface Area

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Temperature

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Alternative reaction routes – catalyst

All based on collision theory

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Maxwell-boltzmann distribution curves

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Total area under the graph must remain constant

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Increasing rate of reaction means that more particles will be past eA on the energy scale

Equilibria

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Dynamic state in a closed system due to reversible reactions

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Le Chatellier’s principle says that if you change the environment, the position of equilibria will change to compensate for the environment

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So increasing the pressure moves towards the less dense side etc etc.

Reactions of Alcohols 

The more methyl groups the C in R-C-OH is bonded to, the more stable the molecule is

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Combustion

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Alcohols and sodium –> H2 and white precipitate (sodium propanoxide)

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Substitution with a halogen produces halogenoalkanes

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This is electrophilic substitution

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Iodo/bromoalkanes need acid catalyst

Oxidisation – goes from orange to green

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Only primary and secondary

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Primary – aldehyde and then carboxylic acid

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Aldehyde + Benedict’s –> Red precipitate

Secondary – ketone

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Ketone + Brady’s reagent –> Yellow/orange precipitate

Halogenoalkanes 

The more reactive the halogen, the less reactive the halogenoalkane

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More methyl groups attached to the C in R-C-X make it more reactive

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Strength of bond can be determined using aqueous silver nitrate solution (hot)

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Faster precipitate = weaker bond

Free radical substitution

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Electron density pulled away from the halogen

Green chemistry bit, with UV light

Reactions summarised:

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+ Hot aqueous alkali –> Alcohol (Nucleophilic substitution)

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+ Concentrated alcoholic alkali (reflux and heat) –> Alkene (Elimination)

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+ Alcoholic ammonia (heat under pressure) –> Amine