As Chemistry Unit 2 Notes

AS Chemistry Unit 2 Notes

Shapes of molecules and Ions

Pairs of electrons will repel each other as far as possible (due to electrostatic repulsion)

Finding the shape:

1. Draw dot and cross

2. Count the number of electron pairs – bond pairs and lone pairs

3. Decide the shape adopted by the electron pairs

4. Look at the number of lone pairs and decide the shape adopted by the atom

5. Draw shape, including bond angles

ELECTRON PAIRS

SHAPE

EXAMPLE

BOND ANGLES AND 3D SHAPE

2 bond pairs

Linear

BeCl2

180o

3 bond pairs

Trigonal Planar

BCl3

120o

4 bond pairs

Tetrahedral

CH4

109.5o

5 bond pairs

Trigonal Bipyramidal

PCl5

6 bond pairs

Octahedral

SF6

90o

Lone pairs repel more than bond pairs because they are attracted to a single nucleus and not

shared by two atoms

Lone pairs reduce bond angles between bonding pairs. Each lone pair reduces predicted

bond angle between bonding electrons by 2.5 degrees.

4 election pairs on the central atom – based on the tetrahedral shape:

4 bond pairs = tetrahedral (e.g. CH4 and NH4+) - 109.5o

ELECTRON PAIRS

SHAPE

EXAMPLE

BOND ANGLES AND 3-D

SHAPE

3 bond pairs and 1 lone pair

Trigonal Pyramidal

NH3

107o

2 bond pairs and 2 lone

pairs

Bent/non-linear

H2O

Organic molecules:

Tetrahedral around carbon if saturated e.g. C3H8 or trigonal planar around carbon if there is

a C=C bond.

In C2H4 , the double bond reduces the bond angle further (its electron rich)

Multiple Bonds:

Count as one bond pair of electrons for purpose of determining the shape.

E.g. CO2 is linear:

Carbon Structures:

Carbon has several allotropes – different molecular structures due to differences in

bonding.

Diamond: Each carbon atom forms 4 identical bonds to neighbouring carbon atoms

giving a tetrahedral arrangement. Because of the strong covalent bonds, diamond

has a VERY high melting temperature, is extremely hard (the hardest known

substance) and cannot conduct electricity – no free e-.

Graphite: Carbon atoms in layers. Within a layer, each carbon atom is bonded to 3

other carbons – the 4th outer e- is delocalised and free to move: conducts electricity.

Layers of graphite are weakly bonded to each other – London forces. Also has a very high

melting point.

Fullerenes: Consist of 32+ carbon atoms. Buckminsterfullerene has 60 carbon

atoms. Ball-shaped molecules. The fourth outer e- is delocalised, so conduct

electricity.

Nanotubes: Fullerenes in the form of tubes. Very small and stiffer than other known

materials. If embedded in polymers they may produce materials with good electrical

conductivity and strength.

Intermediate bonding and bond polarity:

Electronegativity: ability of an atom to attract an electron pair in a covalent bond. Increases

ACROSS the period (Fluorine is the most EN element) and decreases down groups.

Differences in elecronegativity between two elements will result in electrons being pulled

further to one end, and there will be POLARITY in the bond e.g:

If the difference is large enough, electrons will be transferred – IONIC BOND.

Small, highly charged cations (e.g. Al3+) are highly polarizing, and will pull electrons toward

then very strongly, especially from a large anion (e.g. I-), resulting in a covalent bond (AlI3 )

If molecule is SYMMETRICAL, there is no overall polarity. E.g. CCl4. The dipoles

cancel out.

Unsymmetrical molecules containing polar bonds will be polar molecules – describes as

having a permanent dipole.

Polar bonds will deflect a stream of water (because water is polar) e.g. CH3Cl deflects, CCl4

doesn’t.

Intermolecular Forces:

3 types of forces BETWEEN molecules: London forces/Van der Waals (weakest), Permanent

dipole-dipole and Hydrogen bonding (strongest)

1. London/VdW: Found in ALL molecules. Caused by an unequal distribution of

electrons which makes a temporary dipole. This affects surrounding atoms causing

induced dipoles. The net result is a weak attractive force. Everything has London

forces, and the MORE electrons, the STRONGER/LARGER the force.

2. Permanent dipole-dipoles: delta-plus of one molecule is

attracted to delta-minus of another molecule:

3. Hydrogen bonding: The attraction between a hydrogen

attached to Fluorine, Oxygen or Nitrogen on one molecule and an F, O, N atom on

another molecule. e.g.: hydrogen bonding in water:

Trends in physical properties by intermolecular forces:

Alkane boiling points increase with carbon chain length, because the number of electrons

increases, so more London forces.

Branched chain alkanes have LOWER boiling points than straight chains because they can’t

pack as closely together, whereas straight chains can pack together closely (greater surface

area in contact) therefore the IMF forces are greater and they have higher boiling points.

Alcohols have very high boiling points (lower volatility – harder to evaporate) due to strong

hydrogen bonding.

HF has a high boiling point due to hydrogen

bonding. The graph dips down to HCl, HBr and HI,

which all have dipole- dipole interactions but the

number of electrons is increasing, so there are

additional London forces which raise the boiling

points.

Solubiliy:

Affected by bonding, and usually a substance will only dissolve if the strength of the new

bonds formed is the same, or greater than the strength of the bonds that are broken.

Ionic compounds dissolve in polar substances such as water, because the ions are attracted

to the polar molecules and they surround the ions and pull them away from the ionic

lattice. This releases energy known as the hydration enthalpy. This can only happen if the

hydration enthalpy is big enough to overcome the lattice enthalpy. (Hydration vs. Lattice)

Alcohols are soluble in water, because they form hydrogen bonds with it.

Non-polar molecules wont form hydrogen bonds with water, so don’t dissolve in it. E.g.

halogenoalkanes like chlorobutane.

Generally ‘like dissolves like’

Example question: State and explain the solubility of hexane in water

Hexane molecules are held together by London forces. Water molecules are held together by

hydrogen bonds. Hexane can’t make hydrogen bonds with water, so the two liquids do not mix or

dissolve in each other – immiscible.

Redox

Oxidation number: the number of electrons that need to be lost or gained to become a

neutral atom.

Uncombined elements are 0

F is always -1, group 1 are +1, group 2 are +2, oxygen is -2 (except in peroxides H2O2 where

its -1), H is +1 (except metal hydrides where its -1)

Oxidation numbers in a neutral compound add up to zero, and in a polyatomic ion add up to

the charge.

Ionic half equations are used for redox processes – when oxidation and reduction take place

together in a reaction.

If species are reduced, electrons are on the LEFT

If species are oxidised, electrons are on the RIGHT

Half equations are then added together for the full redox equation

E.g. : The overall equation for the oxidation of I- ions by MnO4- ions is obtained from the two

half equations:

MnO4- + 8H+ + 5e- Mn2+ + 4H2O

And

2I- I2 + 2e-

For oxidising agents that contain OXYGEN, e.g. MnO4- , you will need H+ on the LEFT and H2O

on the RIGHT (oxygen can’t swim)

The MnO4- half equation has 5e- but the I- equation has 2e- , so to make them both have the

same number of electrons (so they can cancel out when the equations are added together),

the MnO4- equation has to be multiplied by 2, and the I- equation multiplied by 5, so that

they both have 10e-

They are then added together to give:

2MnO4- + 16H+ + 10I- 2Mn2+ + 8H2O + 5I2

Disproportionation: when one species is both oxidised AND reduced at the same time

e.g.: Cl2 + H2O HCl + HClO

0 -1 +1

The periodic table – Group 2

Have their highest energy electrons in an s sub-shell, hence they are called s-block elements.

Ionization energy (I.E) trends:

Going down the group, there is an extra electron shell compared to the element above, and

the atomic radius is increasing

The outer electrons are increasingly further away from the nucleus; therefore the attractive

force is less.

The extra inner shells shield the outer electrons from the attraction of the nucleus

Therefore, the ionization energies DECREASE down the group (gets easier to remove an e- )

Reactions of group 2 elements with Oxygen, Water and Chlorine:

1. Burn in Oxygen to form solid oxides, often burning with a bright flame

e.g.: 2Mg(s) + O2 (g) 2MgO(s)

2. React with water to form metal hydroxide and hydrogen:

e.g.: Ca(s) + 2H2O (l) Ca (OH) 2 (aq) + H2 (g)

Mg reacts rapidly with steam: Mg(s) + H2O (g) MgO + H2

3. React with chlorine to form solid metal chlorides:

e.g.: Mg(s) + Cl2 (g) MgCl2(s)

Reactions of group 2 OXIDES and HYDROXIDES:

1. Group 2 oxides react with water to form metal hydroxides, which dissolve. They are also

alkaline

e.g.: CaO(s) + H2O (l) Ca (OH) 2 (aq)

2. Group 2 oxides and hydroxides are BASES

They neutralise dilute acids e.g.: HCl or HNO3

Form the corresponding salt and water

e.g.: MgO(s) + 2HCl (aq) MgCl2 (aq) + H2O (l)

CaO(s) + 2HNO3 (aq) Ca (NO3)2 (aq) + H2O (l)

Hydroxides are the same: Ca (OH)2 (aq) + 2HCl(aq) MgCl2 (aq) + H2O (l)

Solubility trends of hydroxides and sulphates:

Generally compounds of group 2 elements that contain singly charged negative ions (e.g.

OH-) INCREASE in solubility down group

Compounds with doubly charged -ve ions (e.g. SO42-) DECREASE in solubility down group.

Reactivity INCREASES

down group, as the I.E

decreases:

Be doesn’t react

Mg with steam

Ca steadily

Sr fairly quick

Ba rapidly

Solubility of

HYDROXIDES

INCREASES down the

group

Mg (OH) 2 Insoluble

Ca (OH) 2

Sr (OH) 2

Ba (OH) 2 Most soluble

Solubility of

SULFATES DECREASES

down the group

MgSO4 Most soluble

CaSO4

SrSO4

BaSO4 Insoluble

Thermal Stability of group 1 and 2 CARBONATES and NITRATES

Thermal decomposition: when a substance decomposes when heated

The more thermally stable a substance is, the more heat it requires to break it down.

The carbonate and nitrate ions are LARGE and can be made UNSTABLE by a cation. The

greater the polarising power of the cation, the greater the distortion and the LESS stable the

anion.

The further down the group, the larger the cations and less distortion caused therefore the

MORE stable the carbonate/nitrate anion. Thermal stability increases down a group.

Group 2 compounds are LESS THERMALLY STABLE than group 1 (more distortion by +2

cation)

Group 1

Carbonates: From sodium carbonate down group 1, the carbonates will NOT DECOMPOSE

on heating – thermally stable.

Nitrates: From sodium nitrate down group 1, the nitrates decompose to form the nitrite

and oxygen

e.g.: KNO3(s) 2KNO2(s) + O2 (g)

Group 2

Carbonates: Lithium and group 2 carbonates decompose to form an oxide and carbon

dioxide

e.g.: CaCO3(s) CaO(s) + CO2 (g)

Li2CO3(s) Li2O(s) + CO2 (g)

Nitrates: Lithium and group 2 nitrates decompose to form the oxide, nitrogen dioxide and

oxygen.

e.g.: Ca(NO3)2 (s) 2CaO (s) + 4NO2 (g) + O2(g)

4LiNO3(s) 2Li2O(s) + 4NO2 (g) + O2(g)

Testing thermal stability of nitrates and carbonates:

1. Nitrates:

How long it takes until a brown gas - NO2 – is produced. It is toxic, so must be done in fume

cupboard

2. Carbonates:

How long it takes for carbon dioxide to be produced – tested using limewater which turns

cloudy.

Potassium

Nitrite

Potassium

Nitrate

Flame tests:

1. Mix small amount of compound with few drops of hydrochloric acid

2. Heat a platinum or nichrome wire in hot flame to clean it.

3. Dip the wire into the compound and hold it in hot flame.

Electrons are being excited to higher energy levels by the heat energy. When the electrons

return to the lower energy levels, they emit energy in the form of visible light.

Flame colours of group 1 and 2 compounds:

Group 1: Lithium – RED Group 2: Magnesium – WHITE

Sodium – YELLOW Calcium – BRICK RED

Potassium – LILAC Strontium – CRIMSON RED

Barium – GREEN

The periodic table – group 7, the HALOGENS

Non-metallic elements, VERY reactive.

Diatomic covalent molecules

OXIDISING agents (they are reduced themselves), and become less oxidising, or reactive

down the group.

Reactions of halogens:

1. Disproportionation with alkalis – NaOH or KOH

COLD alkali to give halide and halate (I) ions: HOT alkali to give halide and halate (V) ions:

X2 + 2NaOH NaXO + NaX + H2O 3X2 + 6NaOH NaXO3 + 5NaX +3H2O

X2(g) + 2OH-(aq) XO-

(aq) + X-(aq) +H2O 3X2 (g) + 6OH-

(aq) XO3- (aq) + 5X- + 3H2O

O.S: 0 +1 -1 0 +5 -1

e.g: I2 + 2NaOH NaIO + NaI + H2O 3Br2 + 6NaOH NaBrO3 + 5NaBr + 3H2O

Halogen Physical state and colour Appearance in water Appearance in hydrocarbon solvent

Fluorine Pale yellow gas N/A N/A

Chlorine Green gas Pale yellow/green solution

Pale yellow/green solution

Bromine Red-brown liquid Red/brown/orange Red/brown/orange

Iodine Grey solid Brown Pink/violet

Sodium

iodate (I)

Sodium

iodide

Sodium

bromate (V)

Sodium

bromate

2. Oxidise metals, non-metals and ions

Metals: e.g. fluorine and chlorine react with iron to form iron (III) halides

Iron is oxidised: 2Fe 2Fe3+ + 6e-

Chlorine is reduced: 3Cl2 + 6e- 6Cl-

Overall equation: 3Cl2(g) + 2Fe(s) 2FeCl3(s)

Non metals: e.g. chlorine reacts with sulphur to form sulphur (I) chloride. Sulphur is oxidised

to +1 and chlorine is reduced to -1)

S8(s) + 4Cl2(g) 4S2Cl2(l)

Ions: e.g. all halogens except iodine (weak oxidising agent) will oxidise iron (II) ions to iron

(III) ions in solution. The solution will change colour from green to orange.

2Fe2+(aq) 2Fe3+

(aq) + 2e-

Reactions of Halides:

1. Potassium halides with concentrated sulphuric acid:

React to give a hydrogen halide.

The trend in strength of the halide ions as reducing agents is: I- > Br- > Cl-

KCl and H2SO4:

KCl(s) + H2SO4(l) KHSO4(s) + HCl(g)

But hydrogen chloride is not a strong enough reducing agent to reduce the sulphuric acid, so

reaction stops there. Misty fumes of hydrogen chloride gas will be seen when it comes into contact

with moisture in air. This is NOT a redox reaction – O.S of halide and sulphur stay the same (-1 and

+6)

KBr and H2SO4:

KBr(s) + H2SO4(l) KHSO4 (s) + HBr(g)

This reaction gives misty fumes of hydrogen bromide gas, and the HBr is strong enough to reduce

the H2SO4 in a redox reaction.

Then this reaction: 2HBr + H2SO4 (l) Br2(g) + SO2(g) + 2H2O(l)

O.S of Br: -1 0 OXIDATION

O.S of S: +6 +4 REDUCTION

KI with H2SO4:

KI(s) + H2SO4(l) KHSO4(S) + HI(g)

2HI + H2SO4(l) I2(s) + SO2(g) + 2H2O(l)

Same first two reactions, but because iodine is a very strong reducing agent, it goes

further, and reduces SO2 to H2S :

6HI(g) + SO2(g) H2S(g) + 3I2(s) + 2H2O(l)

O.S of I: -1 0 OXIDATION

O.S of S: +4 -2 REDUCTION

H2S is a toxic gas, and gives a bad egg smell.

2. Hydrogen Halides with ammonia and water

Hydrogen halides are colourless gases. They are very soluble, and dissolve in water

to make STRONG acids:

HCl (g) H+

(aq) + Cl- (aq) (dissociation)

Hydrogen chloride forms hydrochloric acid; hydrogen bromide forms hydrobromic

acid and so on.

With ammonia: react to form white fumes of the corresponding ammonium halide:

NH3 (g) + HCl(g) NH4Cl(s)

3. Displacement by more reactive halogens

The oxididising strengths of the halogens can be seen in their displacement reaction

with halides.

E.g. Br2(aq) + 2KI(aq) 2 KBr(aq) + I2(aq)

The bromine displaces the iodine ions (it oxidises them) giving iodine I2(aq) and

potassium bromide

A halogen will displace a halide from solution if the halide is below it in the periodic

table

4. Silver nitrate solution, and silver halides solubility in ammonia and reactions with sunlight:

To test for halides in solution:

1. Add dilute nitric acid – this removes ions that could interfere with test and ppt.

2. Add silver nitrate solution (AgNO3(aq) )

A precipitate of the silver halide will form, the reaction is:

Ag+ (aq) + X-

(aq) AgX(s)

The colour of precipitate identifies the halide, and they have different solubility’s in ammonia

solution:

Chloride Cl- : White ppt which dissolves in dilute NH3 (aq) and darkens in sunlight

Bromide Br- : Cream ppt, dissolves in concentrated NH3 (aq) and darkens in sunlight

Iodide I- : Yellow ppt, insoluble in concentrated NH3 (aq) and does NOT darken in

sunlight.

The reaction of silver halides with sunlight (decomposition) is:

2AgBr 2Ag + Br2

Making predictions about fluorine and astatine from trends in group 7:

Number of electrons increases down group, so London forces will increase. Astatine

will be a solid and have the highest boiling temperature.

Electronegativity decreases down group, so astatine will have lowest EN value.

Fluorine will be most oxidising

Kinetics

Reactions only happen when: Particles collide in the correct orientation, and they possess

the activation energy (minimum amount of kinetic energy particles need to react). This is

the collision theory.

Enthalpy profile diagram:

Factors affecting the rate of reaction: concentration, temperature, pressure, surface area

and catalysis.

Factor How it affects rate

Explanation

Concentration (solution)

Pressure (gas)

Increasing conc./pressure increases rate

The particles become more crowded, therefore collide more times which increases the reaction rate.

Surface area (solids)

Increasing surface area increases rate

The smaller the size of reacting particles, the greater the total surface area. Increasing surface area means larger area is exposed for reaction and more collisions.

Temperature

Increasing temperature increases rate

Increasing temperature means the average speed of reacting particles increases, therefore more collisions per second.

Catalyst

Speeds up the reaction

Lower the activation energy by providing an alternative route. If activation energy is lower, more particles will have enough energy to react.

Maxwell-Boltzmann distribution:

Shows distributions of molecular energies in a gas

When temperature is increased, particles will have more kinetic energy and move faster.

This means that more particles will have energies greater than the activation energy and will

react. This changes the shape of the Maxwell Boltzmann distribution curve pushing it to the

right, with a peak lower than the original.

Catalysts:

Increase the rate of a reaction by providing an alternative reaction pathway with a lower

activation energy. It is chemically unchanged at the end of the reaction.

Homogenous catalysts: in the same state as the reactants.

Forms intermediates with the reactants, which the products are then formed from.

The activation energy needed to form the intermediates and the products from the

intermediates is lower than that needed to make the products directly from the reactants.

Only molecules in this

region can react –

molecules have a higher

energy than the

activation energy

Total number of gas

molecules under the

curve

Higher temperature

Lower temperature

Chemical Equilibia

Many reactions do not go to completion because the reaction is reversible

Dynamic equilibrium: When the rates of the forward and reverse reactions are equal. It’s

dynamic because individual molecules react continuously. It is at equilibrium because no net

change occurs (overall concentrations remain constant

Equilibrium can only happen in a CLOSED system.

The effect of conditions on the position of

equilibrium:

Controlled by Le Chatelier’s principle: When a

system at equilibrium is subjected to a

change, it will behave in such a way to

counteract that change.

Temperature is a very important way to

control industrial processes, because it is the

most effective factor (general rule – increase

in 10K doubles the rate of reaction.

Pressure is very expensive to use in

equilibrium processes.

The red-brown gas NO2 exists in equilibrium

with pale yellow N2O4 :

N2O4 2NO2

The forward reaction is endothermic.

If the position of equilibrium shifts to left the

mixture pales

If the position of equilibrium shifts to right the mixture darkens

ORGANICS – Alcohols:

General formula: CnH2n+1OH where the functional group is C-OH

Examples: CH3OH – Methanol – used for fuels and plastics.

CH3CH2OH – Ethanol – fuels, alcoholic drinks

CH3CH2CH2OH – Propan-1-ol

CH3CHOHCH3 - Propan-2- ol

Can be primary secondary or tertiary alcohols:

Reactions of alcohols:

1. Combustion of alcohols:

C2H5OH(l) + 3O2(g) 2CO2(g) +3H2O(g)

2. Reaction with Sodium:

2Na(s) + 2CH3CH2OH (l) 2CH3CH2O-Na+ + H2 (g)

Sodium + ethanol sodium ethoxide + hydrogen

And the longer the hydrocarbon chain, the less reactive with sodium.

3. Substitution reactions to form halogenoalkanes:

Alcohols react with PCl5 (Phosphorus (v) Chloride), releasing hydrogen chloride gas which

forms misty fumes in air

CH3CH2OH (l) + PCl5 CH3CH2Cl (l) + POCl3 (l) + HCl (g)

The OH is swapped for the Cl, and this reaction can be used as a test for an –OH group. The

steamy fumes that are produced turn blue litmus paper red (because HCl dissolves to form

a strong acid)

To make a chloroalkane, just mix a tertiary alcohol (most reactive) and hydrochloric acid

together. This will give an impure chloroalkane which can be purified.

4. Oxidation of alcohols:

Must be familiar with these functional groups:

To oxidise alcohols we use acidified potassium dichromate solution. This is orange in colour and is a mixture of sulphuric acid, H2SO4 and K2Cr2O7. The orange colour is due to the Cr6+ ions in K2Cr2O7. If it oxidises (i.e. the Cr6+ ions become reduced) then the solution turns green.

Cr6+(aq) + 3e- Cr3+

(aq) Orange Green

The results show that only primary and secondary alcohols can be oxidised, and tertiary

alcohols cannot be oxidised, therefore remains orange.

Observations made:

Sodium fizzes,

bubbles form, sodium

disappears, and white

solid product forms

Oxidation of primary alcohols:

A primary alcohol can be oxidised to an aldehyde and then to a carboxylic acid. This is

carried out using an oxidising agent: Mixture of sulphuric acid, H2SO4 (souce of H+) and

potassium/sodium dichromate, K2Cr2O7

To stop oxidising at the aldehyde, you must’ allow the product to distil over’

To get the carboxylic acid, you heat under reflux

Primary

alcohol to

aldehyde

This is the distillation apparatus.

The aldehyde has to be distilled off as it forms

as it can be oxidised further

Distillation evaporates and condenses liquids at

different temperatures. Collect the liquid you

want around its boiling point and discard any

others

Primary alcohol

to carboxylic

acid

When making the carboxylic acid the mixture is refluxed.

Heated strongly with an excess of the acidified potassium or

sodium dichromate, and the alcohol will be completely

oxidised passing through the aldehyde stage to form a

carboxylic acid.

HEAT

Refluxing allows you to heat / boil volatile liquids for a long time. The condenser stops the

volatile liquids evaporating off , because any vaporised compounds are cooled, condense

and drip back down to the reaction mixture

Oxidation of secondary alcohols:

Secondary alcohols are oxidised to ketones ONLY. Do not undergo further oxidation.

This can be done by refluxing the secondary alcohol with acidified sodium/potassium

dichromate.

Summary:

Primary alcohol Aldehyde Carboxylic acid

Secondary alcohol Ketone No reaction

Tertiary alcohol No reaction

Halogenoalkanes

Halogenoalkanes have the general formula CnH2n+1X. X is a halogen.

Can also be primary, secondary and tertiary like alcohols.

When naming, the halogen part is named first (prefix chloro-, bromo-, iodo-) followed by

name of alkane

E.g. CH3Cl = Cloromethane

CH3CH2Br = Bromoethane

If there is more than one halogen di- and tri- are used to indicate the number of halogens

present, e.g. CH2BrCH2Br = 1,2-dibromoethane

Reactions of Halogenoalkanes:

Halogenoalkanes contain polar bonds because the halogen is

more electronegative than the carbon. This leaves a carbon

with a delta + charge, making it open to attack by

nucleophiles.

Nucleophiles: attracted to electron deficient atom, d+ and donate a pair of electrons to

form a new covalent bond

The halogen will be replaced by the nucleophile, which gives a substitution reaction, giving a

new functional group.

1. Halogenoalkanes react with aqueous alkalis to form ALCOHOLS

Aqueous hydroxide ions need to substitute the halogen. Sodium hydroxide NaOH(aq) or

potassium hydroxide KOH(aq) can be used.

The reaction is called hydrolysis and usually carried out under reflux

Hydrolysis: is a reaction with water or aq hydroxide ions that break a chemical compound

into two compounds

Mechanism:

Water can act as a nucleophile too, but it is

a much slower reaction:

First step

Second step

The overall equation with water:

If water with dissolved silver nitrate is used, this can tell us about the reactivities of

halogenoalkanes

When water and an alcohol react, and an alcohol is formed, the silver nitrate will react with

the halide ions when they form giving a silver halide precipitate

The precipitate that forms first indicates which halogenoalkanes hydrolyses first:

Tertiary halogenoalkanes – precipitate forms immediately

Secondary halogenoalkanes – precipitate forms after several seconds

Primary halogenoalkanes – precipitate forms after several minutes

This shows that the reactivity is tertiary 3o > secondary 2o > primary 1o

2. Halogenoalkanes react with alcoholic ammonia to form amines

Ammonia NH3 has a lone pair of electrons, and can therefore act as a nucleophile

Alcoholic ammonia – ammonia dissolved in ethanol.

Heated under reflux

3. Alcoholic alkali to form alkenes

When a halogenoalkane reacts with alcoholic alkali, e.g. potassium hydroxide, KOH in hot

ethanol , an alkene is made

This is an elimination reaction

Heated under reflux

Uses of halogenoalkanes:

Halogenoalkanes are used as fire retardants and refrigerants

Chlorofluorocarbons (CFCs) used to be used in the past because of their unique properties

(non-toxic, non-flammable, unreactive), but it was found that they deplete the ozone layer

in the atmosphere, so are being phased out (see notes later)

Other halogenoalkanes such as hydrofluorocarbons (HFCs) are now used as safer

alternatives.

MECHANISM

Step 1

Step 2

In the second step, and ammonia

molecule removes hydrogen from the NH3

group to form an ammonium ion (NH4+)

This can then react with the Br- ion from

step 1, to form ammonium bromide:

NH4Br

Overall reaction: with ethanol and under

reflux

Mechanisms:

Free radical – species with an unpaired electron

Electrophile – species that accepts a pair of electrons

Nucleophile – species that donates a pair of electrons

Substitution – one species is replaced by another

Addition – joining two or more molecules together to make a larger molecule

Elimination – when a small species is eliminated from a larger molecule

Oxidation – loss of electrons. Also is the gain of oxygen/loss of hydrogen

Reduction – gain of electrons. Also is the loss of oxygen/gain of hydrogen

Hydrolysis – Splitting up using water (usually in form of OH- ions)

Polymerisation – joining together monomers into long carbon-chain polymers.

Redox – any reaction where electrons are transferred between two species

Bond breaking – homolytic and heterolytic:

Homolytic – when the bond breaks evenly, and one electron moves to each atom. This

forms two free radicals as both atoms now have an unpaired electron. The unpaired

electron makes free radicals very reactive.

Heterolytic – when the bond breaks evenly, and both electrons from the shared electron

pair move to one atom. This forms two different species: a positively charged cation – an

electrophile, and a negatively charged anion – a nucleophile

When drawing curly arrows – double headed arrow shows movement of electron pair; single

headed arrow shows movement of single electron.

Should be able to recall these reaction mechanisms from unit 1: Electrophilic addition and free

radical substitution:

Free radical substitution of chlorine in alkanes:

Initiation, propagation, termination Hydrogen bromide to alkenes:

Predicting the type of mechanism:

Polar bonds always break heterolytically

A nucleophile can attack the d+ atom in a polar bond

An electrophile can attack an electron- rich part of a molecule – e.g. the C=C bond in alkenes

All reagents used in AS chemistry (helpful to learn them):

Nucleophiles: OH-(aq), NH3 (alcoholic), PCl5, NaBr/H2SO4 or PBr5, P/I2

Electrophiles: H2, X2, HX

Oxidation [O]: KMnO4 /H+, K 2Cr2O7/H+

Other: KOH in hot ethanol, Cl2 / u.v light, RO –OR/ u.v light (polymerisation)

OZONE

Ozone molecules – O3

The ozone layer is at the edge of the stratosphere

It filters out most of the harmful UV radiation which can damage DNA in cells causing skin

cancer and can also cause eye cataracts.

Ozone is formed when UV radiation from the sun hits oxygen molecules. This forms two free

radicals. The free radicals then combine with other oxygen molecules to form ozone

molecules

O2 + U.V O* + O*

O2 + O* O3

The ozone layer is constantly being replaced, and there is a natural balance between

formation of new ozone and breakdown of ozone molecules : O2 + O* O3

It was discovered that the ozone layer is thinning in places, and a hole in the ozone was

discovered over Antarctica – this means that more harmful UV will reach the earth.

The decrease in ozone concentrations is due to CFCs – chlorofluorocarbons.

Because of their un-reactivity, CFCs don’t decay and reach the upper atmosphere and the

ozone layer, where several reactions happen:

1) CFCs are broken down by UV light, forming chlorine free radicals

CCl3F2 (g) CCl2F*(g) + Cl*(g)

2) The free radicals are catalysts, and react with ozone to form an intermediate – ClO*, and O2

Cl*(g) + O3 (g) O2 (g) + ClO*(g)

ClO*(g) + O3 (g) 2O2 (g) + Cl*(g)

3) The overall reaction is: 2O3 (g) 3O2 (g) (Cl is the catalyst)

H+ = acidified

* = free radical

The Cl free radical is regenerated

and goes on to attack other ozone

molecules. This shows that one

CFC molecule can destroy

thousands of ozone molecules

Nitrogen oxides are produced from car and aircraft engines and thunderstorms. Like

chlorine radicals, NO* also act as catalysts :

NO* + O3 O2 + NO2*

NO2* + O3 2O2 + NO*