Production of materials

Ethylene (ethene)

Although ethylene is a widely used raw material

very little of it is found in either natural gas or

crude oil. Instead it has to be produced from

other hydrocarbons by a process called cracking.

Cracking: process in which large hydrocarbons

are broken down into smaller ones with the help

of heat and/or a catalyst.

Cracking

During the cracking process bonds within the

hydrocarbon molecule are broken. Ethene is

produced in one of two ways:

From crude oil by catalytic cracking of fractions from the distillation column

From natural gas by thermal cracking

One possible reaction involving the cracking of hydrocarbon C15H32 might be

http://www.chemguide.co.uk/organicprops/alkanes/cracking.html

Reactions of ethane and ethene

What do ethane and ethene have in common?

2 carbon atom chain, non-polar (insoluble in water, low melting and boiling points, both undergo combustion with oxygen

How do they differ?

ethene has a double bond and therefore is much more reactive

Reactions of ethene

Like all alkenes, ethene undergoes addition

reactions. Why?

Answer: When the double bond is broken

additional atoms or groups of atoms can be

added – one to each C atom previously linked

by the double bond.

Addition of hydrogen

Addition of hydrogen to ethene results in the formation of ethane.

Ethylene + hydrogen ethane

CH2= CH2 (g) + H2 (g) CH3-CH3 (g)

Alkene + hydrogen alkane

Addition of halogens

Addition of a halogen (eg: Cl, Br) to ethene results in the formation of a haloalkane

ethene + bromine 1,2 dibromoethane

CH2= CH2 (g) + Br2 (l) CH2Br-CH2Br

(g)

Alkene + halogen di-halo-alkane

Addition of hydrogen halides

Addition of a hydrogen halides (eg: HCl) to ethylene also produces haloalkanes.

Ethylene + hydrogen chloroethane

chloride

CH2= CH2 (g) + HCl (g) CH3-CH2Cl(l)Alkene + hydrogen haloalkane

halide

Addition of water

Addition of water (in the presence of an acidic catalyst) to ethylene produces an ethanol

Ethylene + water ethanol

CH2= CH2 (g) + H2O (l) CH3-CH2OH

(l)

Alkene + water alkanol

Reactions of alkanes

1. Combustion reaction: Alkanes burn in air to produce CO2 and H2O

C3H8 (g) + 5O2 (g) 3CO2 (g) + 4H2O (g)

propane + oxygen carbon + water

dioxideAlkane + oxygen carbon + water

dioxide

Reactions of alkanes

1. Substitution reaction: Alkanes react with Cl2, Br2, I2 (halogens) when exposed to ultraviolet light

C6H14 (l) + Br2 (l) C6H13Br(l) + HBr(aq)

hexane + bromine bromohexane

Alkane + halogen haloalkane

Polymerisation

Reaction in which many small molecules

(monomers) combine to form one large

molecule (polymer). There are two main types

of polymerisation reactions:

Addition polymerisation Condensation polymerisation

Addition polymerisation

In the process of addition polymerisation

monomers simply add together without the loss

of any atoms. Basically the double bond opens

out to form single bonds with neighbouring

molecules.

Addition reactions involve unsaturated

hydrocarbons.

Addition polymerisation

Condensation polymerisation

Condensation polymerisation involves a

reaction between two monomers which have

different functional groups. Small molecules

such as water are eliminated during this

reaction.

Carboxylic acid functional group

COOH

Condensation polymerisation

Amine functional group –NH 2

Alcohol functional group -OH

Synthetic polmers

Ethene is the simplest monomer capable of

undergoing addition polymerisation. Some

important synthetic polymers formed from

ethene include: Poly(ethene) (polyethelyne) Poly (vinyl chloride) PVC Poly (styrene) Poly (acylonitrile) PAN Poly (propene) (polypropylene)

Biopolymers

Polymers produced by living organisms are

called biopolymers. Examples include:

Cellulose Starch Proteins Nucleic acids

Alcohols

Alcohols are a family of carbon compounds that contain the hydroxy group (-OH).

Alkanols are a specific group of alcohols where one or more hydrogen atoms in an alkane are replaced by an –OH functional group.

Alkanols are represented by the general formula ROH where R = alkyl group

Ethanol

Alcohols : Nomenclature

Add the suffix ‘ol’ in place of the ‘e’ on the name of the hydrocarbon to which the –OH group is attached.

A number indicates the position of the carbon atom containing the -OH group.

If there are more than one –OH group add the suffixes ‘-diol’, ‘-triol’ and so on.

Ethanol 1,2 ethanediol

Primary alcohol

The carbon atom attached to the _OH group

has two carbon atoms bonded to it.

Secondary alcohol

The carbon atom attached to the _OH group

has two carbon atoms bonded to it.

Tertiary alcohol

The carbon atom attached to the _OH group

has two carbon atoms bonded to it.

Ethanol as a solvent

Ethanol is a good solvent because it is a very

polar molecule. When ethanol and water are mixed

they readily dissolve in each other. This is due to

the polar nature of the O-H bond.

The polar end of the ethanol molecule interacts with other

polar molecules to form dipole-dipole forces or

hydrogen bonds eg: with water

C δ+ O δ-

Hδ+

Ethanol as a solvent

Ethanol and hexane (a non-polar molecule)

readily dissolve in each other.

The non - polar end of the ethanol molecule (the alkyl

chain) forms dispersion forces with other non-polar

molecules. This enables ethanol to act as a solvent for

some non-polar molecules.

Production of Ethanol

1. Hydration of ethanol: industrial ethanol is produced by the acid catalysed addition of water to ethene, represented by the equation:

CH2 = CH2 (g) + H20 (g) CH3 - CH2OH(g)

Production of ethanol

2. Fermentation: process in which glucose is broken down to ethanol and carbon dioxide by the action of enzymes in yeast (these act as catalyst).

C6H12O6 (aq) 2CH3 - CH2OH (aq) +2CO2 (g)

This process is exothermic.

Ethanol as a fuels

The combustion of ethanol is an exothermic

reaction.

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

The amount of heat released can be expressed

as the molar heat of combustion:

‘Heat liberated on complete combustion of one

mole of a substance’

Calorimetry

Calorimetry is a method used to determine heat of combustion. Essentially we measure the change in temperature of measured mass of water heated by the combustion of a measured amount of fuel. This is then used to calculate heat energy release per mole of substance burned.

Molar heat of combustion

1. Find the mass of the fuel burned ??? by weighing the fuel and container before

and after heating

2. Calculate the moles of fuel burned ??? n= m/M

Molar heat of combustion

3. From the rise in water temperature calculate heat produced by combustion of that many moles of fuel???

ΔH = m C ΔT

4. Calculate how much heat could have been produced by one mole of the substance

Oxidation-reduction

Reactions which involve the transfer of electrons are called oxidation-reduction reactions.

OXIDATION = LOSS OF ELECTRONSREDUCTION = GAIN OF ELECTRONS

Zn (s) + 2HCl (aq) ZnCl2 (aq) + H2 (g)

Oxidation: Zn (s) Zn 2+ (aq) + 2e-

Reduction: 2H+(aq) + 2e- H2 (g)

Displacement reactions

Displacement reactions are oxidation-reduction

reactions in which a metal converts the ion of

another metal to the neutral atom.

In these reactions the metal dissolves and the

ions of the other metal are reduced to elemental

metal and deposit out of solution.

Example:

Cu 2+ ions have a greater tendency to gain electrons than Zn 2+ ions.

As a result there is a transfer of electrons from the Zn metal to the Cu (II) ions.

As the reaction proceeds Zn metal dissolves and goes into solution as Zn ions and Cu metal is formed

The activity series

The activity series can be used to predict whether a metal will displace the ions of another metal.

K>Na>Mg>Al>Zn>Fe>Sn>Pb>Cu>Ag>Hg>Pt>Au

The more reactive metal will displace another metal from a solution of its ions.

Oxidation states

In many oxidation-reduction reactions it is not obvious which species has been reduced and which has been oxidised.

To overcome this problem we use a system of assigning oxidation states to atoms to keep track of the number of electrons transferred or shared in oxidation-reduction reactions.

Oxidation state is an arbitrary number assigned according to a set of rules.

Rules for determining oxidation state

1. Uncombined elements have an oxidation state of 0

2. Ions have an oxidation state equal to their charge (eg: Na+ = +1)

3. Oxygen in compounds has a charge of -2 in oxides and -1 in peroxides

4. Hydrogen in compounds has a charge of +1 when combined with non-metals and -1 when combined with metals

5. The oxidation state of a compound or polyatomic ion is the sum of the oxidation states of all its atoms.

Oxidation state

Oxidation involves an increase in

oxidation state.

Half equation Zn (s) Zn 2+ (aq) + 2e-

Oxidation state 0 2+

Reduction involves a decrease in

oxidation state.

Half equation: 2H+(aq) + 2e- H2 (g)

Oxidation state: 1+ 0

Note that the number of electrons lost or gained = change in oxidation state

Electrochemical cells

Redox reactions can be used to generate electricity in a galvanic cell

Example: When zinc metal is placed in CuSO4 solution,

following reaction take place: Zn(s)  +  CuSO4(aq) ZnSO4(aq)  +  Cu(s)

Oxidation:   Zn(s) → Zn+2 + 2e- Reduction:  Cu+2 + 2e- → Cu Overall:      Zn(s) + Cu+2 → Zn+2 + Cu(s)

Galvanic cell

Each of the two parts of the cell is called a half cell.Each half cell is connected by a salt bridge which completes the circuit and allowsions to travel between each half cell.

How does the galvanic cell work?

Zn loses electrons to form Zn ions in solution. The Zn strip dissolves Zn ions travel through the external circuit to

the Cu strip where they are accepted by the Cu ions

The Cu ions are reduced to Cu atoms which deposit on the strip

How does a galvanic cell work?

As reaction continues excess Zn2+ ions build up in the ZnSO4

2- solution

excess negative SO42- ions build up in the

CuSO42- solution

To maintain electrical neutrality in the half cell solutions positive Na+ ions move into the copper half cell from the salt bridge at the same time NO3

- ions move into the zinc half cell

Standard Reduction potentials

Standard reduction potential (EO) is a measure of the relative tendency of a substance to gain one or more electrons compared to the standard hydrogen half cell. The larger the EO value the greater the oxidising power of a substance

e.m.f or voltage of a galvanic cell is the difference in the reduction potentials of the two couples making up the cell.