AS Chemistry Edexcel Unit 1 Core Principles · PDF fileUnit 1.3 Atomic structure & the periodic table (p52 – 77 in text book) 1. Relative atomic mass The relative atomic mass of

AS Chemistry

Edexcel

Unit 1

Core Principles

Unit 1.3 Atomic structure & the periodic table

(p52 – 77 in text book)

1. Relative atomic mass

The relative atomic mass of an element is the average mass of its isotopes compared with the mass of an atom of carbon-12.

2. Relative formula mass

For a compound, the relative formula mass is the sum of the relative atomic masses of all the atoms in the chemical formula.

3. Measuring the mass of an atom

To find the relative atomic mass of an atom, you need to measure its mass and compare it with the mass of an atom of carbon-12. This can’t be done by

weighing.

a. The mass spectrometer

This is the instrument used to find out the mass of a sample.

(Refer to fig 1.3.1 p52 for diagram) There are several different stages to the mass spectrometer;

i. Vaporisation – this is necessary to

ii. Electron gun – this produces due to

ii. Acceleration – achieved by

iii. Velocity selector – to ensure

iv. Deflection – achieved by

vi. Detection – shows the abundance of ions with each different

mass:charge ratio (m/z value)

b. Finding the relative atomic mass from a mass spectrum

(Refer to fig 1.3.2 p53)

Ar of sample = (m/z x % abundance) + (m/z x % abundance) + …

(Refer to green box p54)

c. Uses of mass spectrometry

(Refer to HSW boxes p55-59)

i. Radioactive dating ii. Drug testing

iii. Space exploration

4. The arrangement of electrons in atoms


Over time, ideas about atomic structure changed as technology improved.

a. Thomson’s plum pudding model

(Refer to fig 1.3.13) He knew that atoms were neutral but that they contained electrons with a

negative charge. He suggested the positive ‘pudding’ had negative electrons embedded in the structure.

He also established the mass of an electron as approximately 1/2000 of a

hydrogen atom. Further work proved that the positive protons have a larger mass than that of

electrons.

b. Rutherford’s atom with a nucleus (Refer to fig 1.3.15 p61) Rutherford beamed alpha particles through gold foil. While most of the

particles were detected on the other side, some were deflected. From this he deduced that there was a central positive core (the nucleus) surrounded by

electrons. He was able to use his measurements to calculate the diameter of the nucleus and hence the diameter of an atom.

c. Bohr’s electron shells

If gas is heated, or electrically charged, it gives out light which can then be split through a prism or diffraction grating to form a spectrum known as a line

spectrum or emission spectrum. This is always identical for any given atom.

(Refer to fig 1.3.16 p61) Bohr came up with the idea that electrons were in shells around the nucleus.

Moving out from the nucleus successive shells got closer together, in the same way that lines from the emission spectra for hydrogen did. Bohr’s new

model matched exactly the line spectrum for hydrogen. Unfortunately, this only works for hydrogen. Heisenberg & Schrödinger

continued this work, beginning quantum mechanics.

5. Energy levels & electron shells

The absorption and emission of light by an atom can be explained by electron movements between different fixed energy levels i.e. electron shells.

Light is a form of energy. When atoms receive this light energy absorption takes place and the electrons become ‘excited’ and are ‘promoted’ to a higher

energy level (shell). Conversely, emission occurs when electrons fall back to a lower energy level (shell).


Each energy level (shell) is given a principal quantum number, n. So the ground state for hydrogen is n = 1, since hydrogen has its electron on the

first shell (energy level). Each successive energy level rises in number by one each time (2,3,4,5….to infinity).

6. Ionisation energies

Ionisation is the complete removal of an electron from an atom, which is an endothermic process since energy must be used to overcome the forces of

attraction between the electron and the nucleus. Comparing the ionisation energies of different atoms enables a picture to be

built up of the electronic structure for each atom.

An atom in its ground state is at its lowest energy level. The energy required to remove the first electron is called the first ionisation energy, the energy

required to remove the second electron is called the second ionisation energy and so on. To calculate the total energy of removing two electrons you must add the first and second ionisation energies together.

7. Subshells

Each electron shell may contain several different subshells. These are described by letters: s, p, d, f, g. The following subshells are available in

each shell:

Shell Subshell(s)

1 1s

2 2s, 2p

3 3s, 3p, 3d

4 4s, 4p, 4d, 4f

Within a shell, the subshells have different energy levels, with electrons in the

lowest energy subshells being closest to the nucleus:

s (lowest energy) < p < d

Each type of subshell contains one or more orbitals.

First shell (n=1)

Second shell (n=2)

Third shell (n=3)

Fourth shell (n=4)

Subshells

s s p s p d s p d f

Number of orbitals

1 1 3 1 3 5 1 3 5 7

8. Electron spin


Electrons in atoms behave like tiny magnets. A moving charge can create a magnetic field. Electrons can spin either clockwise or anticlockwise. Two electrons in the same orbital cannon have the same spin. This means that

each orbital can have a maximum of 2 electrons, having opposite spin.

Subshell No of orbitals Max No of electrons in subshell

s 1 2

p 3 6

d 5 10

f 7 14

9. Electron configurations (filling the orbitals)

Each orbital is filled in order, lowest energy orbital first. This means a shell is not always completely filled before electrons start filling the next shell e.g. a

4s orbital will fill before a 3d orbital. The order of filling is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p…


Hund’s rule states that electrons in the 2p subshells are placed in different orbitals, rather than placing them in the same orbital with opposite spin.

They ‘spread out’ to maximise the number of unpaired electrons. Once each orbital has had a single unpaired electron placed in it, then the electrons begin to pair up until the subshell is completely filled. (This is true of all

subshells with multiple orbitals.)

(Refer to green box p65)

There are a few exceptions to this in order to reach a more stable energy state e.g.

Copper

Chromium

10. Electron configurations for ions

For an ion, you simply add or subtract the number of electrons to form the ion required.

11. Shorthand notations

The previous noble gas can be used to avoid writing out long electron configurations.

(Refer to HSW box p66)

12. Electron density maps

Modern scientists regard electrons as waves rather than particles. An electron behaves as though it is spread out in an electron cloud surrounding

the nucleus. Quantum mechanics calculations provide electron density maps which plot the most likely positions of the electrons. An orbital is the region

where the probability of finding an electron is greatest. (Refer to fig 1.3.22 p66) From these maps, scientists have determined the shapes and orientations of

orbitals within subshells. (Refer to fig 1.3.23 p67)

13. Structure of the periodic table

Elements are arranged in order of atomic number. Each column is called a

group, while each row is called a period. All elements in the same period have the same number of electron shells i.e.

the same quantum number (e.g. n=3 for period 3).

The elements in each group or period show similar characteristics in either chemical or physical behaviour.

14. The development of the periodic table

The periodic table took on several different forms before the current model was settled upon.


a. Early ideas about atomic mass

1799 – Joseph Proust showed that mass proportions for each element remained constant in compounds

1800 – John Dalton made the initial proposal that elements were made up of atoms. Unfortunately, Dalton wasn’t very good at working out the atomic masses

1828 – Jons Jacob Berzelius published an accurate list of atomic masses, but thanks to Dalton’s earlier mistakes, no one took any interest

b. Doberiener’s triads

1829 – Johann Doberiener showed that many of the then known elements could be arranged in groups of three i.e. triads. This was the first time

elements had been grouped together

1862 – Alexandre-Emile Beguyer de Chancourtois showed similarities between every eighth element. However, the diagram wasn’t published with his work

causing it to be ignored.

c. Newlands’ law of octaves

1863 - John Newlands arranged elements in groups of eight by atomic mass. His work was dismissed since it contained some major flaws. He’d assumed all elements had been discovered so didn’t account for any gaps.

d. Mendeleev’s periodic table

1869 – Dmitri Mendeleev and Julius Meyer published clear presentations. Meyer plotted physical properties against atomic mass to produce curves

demonstrating periodic relationships. However, Mendeleev published first producing a table with gaps in it for elements yet to be discovered. He

predicted the properties of the missing elements and when more were discovered their properties matched his predictions. Consequently,

Mendeleev is credited with producing the modern periodic table.

15. The blocks of the periodic table

The periodic table can be divided into blocks determined by the electron

orbitals which are being filled.

a. The s-block elements This is made up of groups 1 and 2 since their outer electrons are in an s

orbital. These elements are reactive metals since the s electrons are easily lost. They have low melting & boiling points, and lower densities than other

metals but can still conduct electricity.

Hydrogen and helium are also s-block elements using these criteria. However, both are non-metals and so scientists treat them as a separate group.

b. The d-block elements

These are called transition metals, although some of the outer d-block elements don’t share the same characteristics, e.g. zinc. They are much less

reactive since the d orbitals are being filled after the outermost s orbital.

They tend to follow the general properties of metals, e.g. ductile, malleable, sonorous, conductors etc. Some also make very good catalysts.

c. The f-block elements

The top row (lanthanides) are all similar metals. The bottom row (actinides) are radioactive. Some have been synthesised and are very unstable.

d. The p-block elements

This block contains all the non-metals (except hydrogen & helium), along with

the metalloids (semi-metals) and some metals. The p-block metals are relatively unreactive and doesn’t exhibit strong metal

character.

The metalloids occur in a diagonal block. They generally behave as non-metals but can conduct electricity e.g. silicon.

The non-metals form covalent bonds with each other and ionic bonds with metals. Most non-metals exist as small molecules.

The noble gases are extremely unreactive. However, some compounds have

been formed. (Refer to HSW box p73)

16. Trends in the periodic table

Physical and chemical trends occur both across a period and down a group

within the periodic table. (Any specific examples are based on period 3)

a. Atomic radius

In period 3, all the outer electrons are in the third energy level: n=3. Going

across the period, a proton is added to the nucleus of each element. This increases the nuclear charge and therefore increases the attraction between the nucleus and the electrons. The electrons are pulled closer to the nucleus,

thereby decreasing the size of the atomic radius.

Going down a group

b. Ionic radius

For the metal ions – the number of electrons decreases, while the proton number increases across the period. This means the electrons are pulled

much closer to the nucleus, decreasing the ionic radius. For the non-metal ions – the number of electrons increases as the number of

protons increases. The nucleus is less able to pull the electrons towards it, so the ionic radius increases.

Going down a group

c. Periodic trends in ionisation energy

(Refer to fig 1.3.33 p76) The general trend shows an increase across the period. As you move across

the period the nucleus of the atoms contains more protons and therefore a higher charge density. This results in the electrons being attracted further to

the nucleus across the group making the ionisation energy increase.

d. Melting & boiling points (Refer to fig 1.3.34 p77) The general trend shows a decrease across the period but there is a spike for

silicon. This can be explained by the bonding of the elements. Na Al Increases because the atoms form ions which are surrounded by a

sea of delocalised electrons. Across the period the charge density of the ions

increases causing an increase in electrostatic attractions within the lattice structure.

Silicon forms a macromolecular structure with strong covalent bonds that need to be broken to change state.

The other elements form simple covalent molecules. The stronger the

intermolecular forces (generally stronger for larger molecules) the higher the melting & boiling points.

Unit 1.4 Bonding


1. What is a chemical bond?

A bond is the force holding together atoms together. The physical and chemical properties of a molecule depend on the type of bond holding together the atoms.

2. Ionic bonding

These are formed by metals and non-metals bonding together via electrostatic forces of attraction. Consequently they are also known as

electrovalent bonds. The ions then become arranged in a giant lattice structure. Within the lattice the attractive forces between oppositely charged

ions are maximised whilst the repulsive forces between like charges are minimised. The forces within the lattice therefore act equally in all directions.


Ions are formed when electrons are either lost or gained. The movement of electrons seeks to obey the octet rule;

‘When elements react, they tend to do so in a way that results in an outer shell containing eight electrons.’

Example

Na+ 1s2 2s2 2p6

This then becomes isoelectronic with Ne, a noble gas (Ne 1s2 2s2 2p6) i.e. they share the same number and arrangement of electrons.

Further example

Cl-

Ar So, Na loses an electron to obey the octet rule, whilst Cl gains an electron to

do so.

When Na and Cl react together to form NaCl, the ionic substance produced has different properties to the constituent elements due to the lattice

structure.

3. Dot and cross diagrams

These show the transfer of electrons between atoms in order to form ionic compounds. Only the outer electrons are shown in order to simplify the

diagrams (inner electrons do not get directly involved in the bonding process).

The dots and crosses distinguish between electrons of the cations and anions. (Cations are positive ions, while anions are negative ions.)


4. Trends in ionic radii The ionic radius is the radius of an ion in a crystal. The radius of a cation is

smaller than the element’s atomic radius, whilst the radius of an anion is larger than the element’s atomic radius.


5. Types of lattice structure

X-ray diffraction can produce electron density maps of lattice arrangements. These show that the exact arrangement of ions in a lattice can vary

depending on the relative sizes of the different ions present.

Sodium chloride exhibits a face-centred cubic structure. Each ion has 6 neighbouring ions, so has a coordination number of 6. This is a common ionic structure.

(Refer to fig 1.4.5a p81) Caesium chloride exhibits a body-centred cubic structure since caesium is a

larger cation than sodium. More chloride ions can fit around each cation, so the coordination number in this instance is 8.

(Refer to fig 1.4.5b p81)

6. Evidence for the existence of ions (Refer to HSW boxes p82-83)

Physical properties

i. Electrical conductivity

Ionic solids do not conduct electricity. This is because they are in a rigid lattice structure.

However, ionic compounds will conduct electricity when either molten or in solution since the ions are free to move.


ii. Strength

Unlike metals, ionic compounds are not malleable or ductile. The lattice structure for ions is a regular pattern but quite rigid. A sufficient force will

cause the layers to split apart resulting in a brittle nature.

iii. Electrolysis Electrolysis experiments show that the ions can be separated using an

electrical current.


iv. Electron density maps Measurements from electron density maps are able to pinpoint the positions

of ions, along with their size and charge density.


7. Lattice energy The formation of an ionic lattice involves a release in energy so has a

negative enthalpy. By comparison, bond formation requires energy so these have positive enthalpies.

Lattice energy can be calculated using the Born-Haber cycle.


a. What affects lattice energies?

Lattice energy is affected by both the size and charge of ions. The lattice energy becomes less negative as the size of the ions increases.

The lattice energy becomes more negative as the charge on the ions

increases.

The electrostatic force of attraction can be calculated using Coulomb’s law. (Refer to equation p85)

b. Predicting stability

The more exothermic a compounds lattice energy is, the more stable it is.

Theoretical lattice energies can be used to make predictions about the stability of less familiar elements.


c. Polarisation in ionic bonds

Ionic bonds can be distorted by the attraction of the positive cation to the negative outer electrons of the anion. If the distortion is great enough, the bond can begin to exhibit some covalent character.

The polarising power of the cation depends on the charge density which is

linked to the charge and the ionic radius. Cations with a small ionic radius have a greater polarising power e.g. Al3+.

The larger an anion is, the more likely it is to be polarised.


When theoretical and experimental lattice energies are compared, some do not have close agreement for the same compound. If there is a small

difference in electronegativities there is more chance of electron sharing and therefore an increase in covalent character. The theoretical model assumes that the charge is evenly distributed across an ion, and that all ions are

spherical and separate.

8. Covalent bonding

Covalent bonds are generally formed between non-metals. The bonds are formed through the sharing of electrons.

The covalent bond is a balance between the attractive force pulling the two nuclei together (due to the electron density between the nuclei) and the

repulsion between two positively charged nuclei. This balanced distance is the bond length. The amount of energy required to form the bond is the bond enthalpy.

9. Dot and cross diagrams

These diagrams show how the outer electrons are shared between atoms, but

they do not show the actual positions of bonding atoms in the molecule. (Refer to fig 1.4.16 p89)

10. Dative covalent bonds

Sometimes both electrons in a covalent bond come from the same atom. This is known as a dative covalent bond or a coordinate bond.

Example 1 The ammonium ion

Example 2 Aluminium chloride dimers

These bonds are often found in oxides, e.g. carbon monoxide.

11. Evidence for the nature of covalent bonds


a. Molecule sizes These are strong bonds, with most covalent compounds existing as simple

molecules. These have low melting and boiling points.

However, there are some exceptions e.g. carbon allotropes, silicon (IV) oxide SiO2. These have high melting and boiling points.

b. Electron density maps and the shapes of molecules

Electron density maps show that covalent bonds are directional towards areas of electronegativity. Lone pairs of electrons (pairs of unshared electrons) also

form dense areas of electronegativity. Each area of electronegativity repels the others, although lone pairs tend to have a stronger repulsion. This gives covalent molecules a very definite shape within a 3D spatial arrangement.


Some of the more common molecule shapes are;

Shape of

molecule

Example Bond angle No of lone pairs

Linear

Trigonal planar

Tetrahedral

V-shaped

NH3

12. Metallic bonding

Metals are good conductors of heat and electricity. There is a strong attraction between the positive ions and the ‘sea of delocalised’ electrons.


a. Properties of metals


i. Conducting electricity

All metals conduct electricity well. The electrons are able to flow through the structure to ‘carry’ the current.

ii. High thermal conductivity

The electrons can transmit kinetic energy rapidly through the lattice structure. The movement of electrons with a high kinetic energy is random so the energy can be transferred to cooler regions of the metal.

iii. High melting & boiling temperatures

Forces between the metal ions are large, so require a lot of energy to overcome them. The delocalised electrons also act as a ‘glue’ to hold the

structure together.

iv. Malleable & ductile They can be hammered into shape and drawn out into wires. The lattice

structure, unlike in ionic compounds, has some flexibility. The ions are able to move their positions slightly with the movement of the delocalised

electrons.

b. Using metals

Copper was used for the first time about 10 000 years ago. In the Middle

East steel was produced about 4000 years ago, but it was about 2500 years ago when this began in Europe.

The properties of metals make them useful and practical for a variety of different uses.


Unit 1.5 Introductory Organic Chemistry

(p94-107 in text book)

1. What is organic chemistry?

Organic chemistry is considered to be the study of carbon chemistry. Carbon forms a vast number of compounds due to its ability to form single, double and triple bonds.

(Refer to HSW box p94) 1807 – Jons Jacob Berzelius observed that chemicals could be divided into two groups depending on their behaviour. He classified organic molecules as

those that would burn or char on heating. At this time, scientists believed that organic molecules could only come from living things and that they

couldn’t be synthesised (produced artificially).

1828 – Friedrich Wohler demonstrated synthesised ammonium cyanate. He demonstrated that this was the same as the urea found in urine.

1846 – Christian Schonbien accidentally combined nitric and sulphuric acids with cellulose. This produced nitrocellulose, an unstable explosive.

2. The vast range of organic compounds

Carbon forms around 7 million different compounds. Some of these are complex molecules that make up living cells, enzymes, plastics etc.

(Refer to HSW box p95) Many drugs are organic compounds, and do actually come from plant or

animal extracts. E.g. aspirin

3. The difference between hazard and risk

Many organic substances require special handling.

a. Hazard

This is essentially the problem associated with handling the chemical e.g. toxic, corrosive etc

b. Risk

This is the chance that it will actually cause harm. It is dependant upon precautions such as the volumes being used, the expertise of the person handling the chemical, the safety equipment available.

c. Ways of reducing risk

(Refer to p97 & 98)

i. Working on a smaller scale – it is easier to contain the reaction e.g. in a fume cupboard

ii. Taking specific precautions or using alternative techniques – the lowest concentration for a given chemical that works in the

reaction can be used rather than more harmful concentrations iii. Careful use of safety measures – using equipment such as

goggles and fume cupboards iv. Changing the conditions under which a reaction takes place –

this could be to lower the temperature to reduce excessive

fumes. The equilibrium position may be changed but the same products are given off

v. Using less hazardous substances – in some reactions alternative substances can be used to reduce the hazards involved.

CLEAPPS provides a list of suitable alternatives for many reactions

Pesticides are chemicals whose use must be monitored and controlled due to the potential hazards associated with their use. DDT is one such chemical

that resulted in toxic build up in food chains. (Refer to HSW box p98 & 99)

4. The properties of the carbon atom

Carbon forms 4 covalent bonds. Carbon atoms are unique in their ability to form covalent bonds with carbon and other non-metal atoms at the same

time. It can also form single, double or triple bonds with itself. Consequently, it can form a large array of different molecules.

5. Classifying organic compounds

There are millions of different organic molecules that can be classified in a number of different ways.

a. Carbon chain arrangement

i. Aliphatic molecules have straight or branched chain carbon skeletons e.g. alkanes

ii. Alicyclic molecules consist of closed ring structures e.g. cyclobutane iii. Arenes are all derived from the benzene molecule (covered in A2)

b. Functional groups

There are a variety of functional groups that can be formed within organic molecules. Each homologous series has a general formula. The molecules

within the series will have similar chemical properties due to the functional group but will have varying physical properties e.g. boiling points.

(Refer to table 1.5.2 p102)

Homologous series

General formula

Functional group

Example

6. Representing organic compounds

There are several different ways of representing organic molecules, depending on the level of detail required.

(Refer to p103) E.g. propan-1-ol

Empirical formula

Molecular formula

Structural formula

Displayed formula

Skeletal formula

7. Nomenclature (naming molecules)

The IUPAC (International Union of Pure and Applied Chemistry) system is used in naming organic molecules.

a. Prefix

The prefix of an aliphatic compound relates to the number of carbon atoms that make up the longest chain within the molecule.


Prefix

Number of carbon atoms in the main carbon chain

Meth-

1

b. Suffix

The suffix of the name relates directly to the main functional group found in the molecule.


Suffix

Number of carbon atoms in the main carbon chain

-ane

alkane

c. Branching

Organic molecules often have branches. The alkyl groups that form the side

chains have the general formula CnH2n+1 E.g.

(Refer to p105)

Alkyl group

Formula

methyl

CH3-

ethyl

propyl

butyl

pentyl

hexyl

The position of the side chain is indicated by the lowest possible number e.g.

3-ethylhexane

If more than one side chain is present, these are named in alphabetical order

e.g. 3-ethyl-2-methylpentane

If the side chains are on the same carbon atom the numerical position is repeated e.g. 2,2-dimethylpentane

d. Isomerism

Isomers share the same molecular formula but have a different arrangement of atoms. There are several different types of isomerism, some of which won’t be covered in this unit.

i. Structural isomerism

There are 3 types of structural isomers.

Type of structural isomer

Examples

Chain

Butane 2-methylpropane

Positional

Propan-1-ol Propan-2-ol

Functional group

Propanal Propanone

There are some pitfalls in drawing isomers which must be avoided; kink in the chain

rotated molecule

(Refer to p107)

ii. Stereoisomerism

These are molecules that are non-super imposable mirror images of each other. This arises when the 3D arrangement of bonds in the molecule allow

different possible orientations of the atoms.

Unit 1.6 The Alkanes – A Family of Saturated Hydrocarbons


1. Hydrocarbons

Hydrocarbons are molecules made from hydrogen and carbon only. Several different homologous series make up the hydrocarbons. They are all insoluble in water and given sufficient oxygen, will combust to produce

carbon dioxide and water. The 3 families are alkanes, alkenes and alkynes.

2. General properties of the alkanes These are saturated hydrocarbons, so have the maximum amount of

hydrogens bonded to the carbon atoms. The general formula is CnH2n+2 and these molecules exist as both straight chained and branched structures.

(Refer to fig 1.6.1 p108) The boiling points of alkanes increase as the molecular formula increases i.e. as more carbon atoms are added.

However, straight chain molecules have higher boiling points than an isomer

that is branched. The straight chain molecules are able to exist closer to each other and as a result have stronger intermolecular forces of attraction to

overcome.

3. Where are alkanes obtained from?

(Refer to fig 1.6.2 p109) The most important source of alkanes is the fossil fuels. Oil was formed

millions of years ago in the seas. Tiny plants and animals fell to the bottom and over time layers formed above them. Compression in the absence of oxygen caused the formation of crude oil and natural gas. The oil and gas

become trapped under non-porous rock which must be drilled through to access the fuel.

Coal was formed in a similar way from plant remains (mainly large ferns) on land.

4. The economic importance of crude oil

(Refer to HSW box p110) Modern life depends on crude oil which provides fuel for transport and generating electricity. It also provides many of the raw materials needed in

the chemical industry. However, it is a finite resource and most deposits are found in politically sensitive areas of the world.

Until the 1970s oil was regarded as a cheap resource. Then the Middle

Eastern countries supplying the oil decided to produce less and increase their prices. Since then, oil prices have varied greatly (even on a daily basis).

However, the general trend has been an increase. This has a knock on effect on users of oil, e.g. transport costs have increased.

5. Using crude oil

The crude oil extracted from the ground is of little use without processing. It is a mixture of many different hydrocarbons. It undergoes primary distillation (fractional distillation on an industrial scale) to separate the different

fractions. It can then undergo further fractional distillation to give a pure yield of a particular alkane.


Fraction Percentage Length of carbon chain

Use(s)

Refinery gas

Gasoline

Kerosene

Diesel oil / Gas oil

Residue

6. Different fractions in demand

Different reserves of crude oil have the different fractions in varying quantities. The lighter fractions are in much higher demand than the heavier

fractions, which can make up to 50% of the crude oil mixture.


Cracking is used to break down the longer carbon chains into the alkanes found in the lighter fractions along with some short chain alkenes such as

ethane, which are also extremely useful.

Cracking requires extremely high temperatures, so catalysts are often used. This process is thereby known as catalytic cracking. The catalysts used are

often crystalline aluminosilicates (zeolites). (Refer to fig 1.6.6 p112)

7. Knocking and the need for catalytic reforming


The alkanes in petrol are generally 5-10 carbon atoms in length. The power is produced by the explosive combustion in cylinders.

The smooth running of the car depends on the explosion happening at exactly

the right time. If the explosion happens too early ‘knocking’ occurs which produces a loss of power. This is usually the case when the petrol contains a large proportion of straight chain alkanes since they easily ignite. Petrol with

a high proportion of branched chain alkanes are more efficient fuels since they take longer to ignite.

Fuels are given an octane rating to indicate the ratio of straight chain to

branched chain molecules. A fully straight chain fuel has a rating of 0, while a fully branched chain fuel has a rating of 100.

Fuel pumps show the octane rating by the Research Octane Number (RON). Ordinary unleaded is 95 while super unleaded is 98.

Most modern cars have high-performance engines so require fuels with a high

octane rating. Tetraethyl lead(IV) used to be used to prevent knocking but resulted in lead pollution. Nowadays, catalytic reforming is used to produce branched chain molecules from straight chain molecules in a similar process

to catalytic cracking.

8. The chemical properties of the alkanes Apart from combustion, the alkanes are very unreactive. This makes them

very useful. They are non-corrosive with metals making them good lubricating oils. They are also harmless to the skin e.g. petroleum jelly.

At room temperature the alkanes are unaffected by concentrated acids and

alkalis. They are not affected by oxidising agents or reactive metals. The bonds in alkanes involve a very even sharing of electrons since the electronegativities of carbon and hydrogen are very close.

Almost all the reactions with alkanes occur due to the formation of free radicals, which contain an unpaired electron. They have high activation

energies, but once overcome the reaction proceeds quickly in the gas phase.

9. Breaking bonds

This is known as bond fission. Since a covalent bond involves the sharing of 2 electrons, there are 2 ways in which the electrons can be shared out when a bond is broken;

a. Homolytic fission

This involves an equal sharing of the once bonding electrons. As a result, each atom in the bond receives an unpaired electron to form a free radical.

These free radicals are extremely reactive. The reaction of one free radical with a substance usually results in the formation of another free radical.

(Refer to fig 1.6.8 p114) Free radicals in the body are now thought to be carcinogens. The number of free radicals in the body can be reduced by cutting down on the intake of

food high in free radicals – burnt toast and charred food from the BBQ for example. Fresh fruit and vegetables rich in vitamins A, C and E help the

enzyme superoxide dismutase to deactivate free radicals to prevent any damage they may cause.


Free radicals are also linked to the aging process. The theory is that they attack the cross links between collagen fibres in the skin causing it to appear

less flexible.

They also attack the DNA in nuclei in body cells which is thought to be the cause of cancer.

Antioxidants are viewed to ‘mop up’ the free radicals in the body. Vitamin E is a well known example of a natural free radical inhibitor.


b. Heterolytic fission

This involves an unequal sharing of the once bonding electrons. This results in the formation of ions.


10. The reactions of alkanes

There are just 2 common types of reactions that alkanes undergo

a. Cracking

When heated to high temperatures in the absence of air alkanes will undergo cracking. This is a thermal decomposition reaction to form smaller molecules.

b. Combustion

When heated in a plentiful supply of air, combustion will take place. This is a highly exothermic reaction. Gaseous alkanes will burn completely but solid and liquid fuels are difficult to ignite.

This reaction is highly important since it is used to generate electricity, fuel

fires in the home, provide central heating, cooking and transport.

(Refer to fig 1.6.12 p116) Methane, propane and butane are all commonly used as fuels. Where

possible, natural gas is piped into homes. When this can’t happen, canisters of propane can be used. Propane can be readily liquefied making it easy to

store and transport. It won’t burn until it returns to the gaseous state making it safe to use.

Incomplete combustion can produce carbon monoxide and has resulted in fatalities. Regular maintenance of gas appliances is required to reduce the

risk of CO poisoning. There are strict regulations imposed on landlords to ensure the safety of tenants.


c. Reaction with chlorine

This reaction is known as free radical substitution, where chlorine radicals substitute for hydrogen atoms on an alkane.

The reaction requires a large input of energy in the form of UV light. This provides the energy to break the bond in chlorine molecules and is known as

initiation. It is an example of homolytic fission.

In propagation the chlorine radical reacts with the alkane. At the end, another free radical is generated. This process can be repeated hundreds of

times as an explosive chain reaction.

(Refer to fig 1.6.14 p118) In termination, 2 radicals combine in a highly exothermic reaction. This

happens every few thousand reactions.

(Refer to fig 1.6.15 p118) Provided there is an excess of chlorine, further substitutions can take place.

Two of these compounds are well known. Trichloromethane (chloroform) CHCl3 was one of the first anaesthetics used. Tetrachloromethane CCl4 was

widely used as a solvent until its carcinogenic properties were recognised.

11. Ethical issues

(Refer to HSW boxes p120 – 123)

a. Cars and society

One of the major uses of alkanes is as fuel for transport. Cars are one of the major causes of pollution on the planet. Carbon monoxide is formed in the

engine which is a greenhouse gas as well as being highly toxic. Oxides of nitrogen and sulphur are also produced which cause acid rain and smog.

Catalytic converters convert CO and NOx into less harmful gases but cannot remove the CO2.

b. The greenhouse effect

Greenhouse gases include water vapour, CO2 and methane, 2 of which are products of hydrocarbon combustion. These gases cause the temperature on

Earth to rise due to re-radiation of energy from the Sun.

c. Developing new fuels

Biofuels such as ethanol, along with hydrogen powered and solar powered vehicles have been looked at as alternatives. However, the amount of land

needed to grow the plants for biofuels is vast.

d. Battery-powered vehicles

Electricity is an alternative, but this is often produced through the combustion of fossil fuels. The technology is still in its early stages and the size of the battery required is large and heavy providing far less energy than petrol.

Recharging can then take up to 12 hours.

e. The hydrogen cell Electrolysis of hydrogen takes place in a cell to form water and heat. The

hydrogen cell motorbike can travel at around 50mph for 100 miles at a cost of £2. However, this technology is in its early stages. Petrol stations aren’t yet

equipped to offer refuelling of hydrogen, even though it is relatively safe.

Unit 1.7 The Alkenes – A Family of Unsaturated Hydrocarbons


1. The carbon – carbon double bond

These are unsaturated hydrocarbons due to the existence of the double bond, which makes them more reactive. There general formula is CnH2n

In a single bond (sigma ơ bond), the electron cloud is symmetrical about the central axis of the molecule.

In a double bond not only does the ơ bond exist, but so does a pi π bond. The electron density is concentrated in 2 regions on either side of the bond

axis. Consequently, there is no rotation around the double bond.


The π bond is reactive towards electrophiles, providing an electron pair to bond with electron seeking species. It also results in geometric isomerism due to the lack of rotation around the double bond.

2. Ethene

(Refer to HSW box p125) Ethene is a plant hormone used in ripening fruit. It is used on a commercial basis to ripen fruit ready for supermarket shelves. However, this could be

having adverse affects on the planet due to the many miles travelled in transporting some fruit across the globe.

3. Geometric isomerism

These have different physical properties e.g. boiling point, but identical chemical properties. They are a form of stereoisomers since there are 2

possible arrangements of the species on the double bond.

The traditional way of naming these is cis-trans isomerism e.g. Cis-but-2-ene Tran-but-2-ene

However, this method of naming has its limitations. The IUPAC system is now the E-Z isomerism system. This is capable of working for all geometric

isomers.

Groups around the double bond are ranked on their atomic number. The atom with the highest number has the highest priority. If the two groups

with higher priority are on the same side they are zusammen, the Z-isomer. If the higher priority groups are on opposite sides they are entgagen, the E-isomer.

E.g. 1-bromo-2-chloro-2-fluoro-1-iodoethene

Z-isomer E-isomer

The body is sensitive to geometric isomers. This has been studied in cooked tomatoes.


4. Reactions of the alkenes

Alkenes are more reactive than alkanes due to the presence of the double bond. The high electron density of the double bond means they are attacked by electrophiles. An electrophile is an electron deficient species that can form

a covalent bond e.g. H+ ion

a. Electrophilic addition reactions

i. Reaction with hydrogen Alkenes can be reduced to alkanes through addition of hydrogen at around

200°C with Ni as a catalyst.

ii. Reaction with the halogens

A halogen molecule can be added across the double bond to produce a dihalogenoalkane e.g. ethene reacting with chlorine. This is possible due to

the temporary dipole in the halogen molecules.

iii. Testing for alkenes with bromine water

The major product is 2-bromoethanol but some 1,2-dibromoethane is also formed. Bromine water is decolourised in a positive test.

iv. Reaction with hydrogen halides

This reaction proceeds rapidly at room temperature due to the permanent

dipole in the hydrogen halide . The resulting compound is a halogenoalkane. In asymmetric alkenes, there are 2 possible products, but not in equal

proportions. The major product can be predicted using Markovnikov’s rule:

When HX adds across an asymmetric double bond, the major product formed is the molecule in which hydrogen adds to the carbon atom in the double bond with the greater number of hydrogen atoms already attached to it.

v. Reaction with acidified potassium manganate(VII)

This is an oxidation reaction. The potassium manganate(VII) is decolourised in the production of a diol. This reaction can also be used to distinguish

between alkenes and alkanes.

It is a useful product, but is produced industrially from epoxyethane.

vi. Reaction mechanism

This is an electrophilic addition mechanism. A carbocation (carbonium ion) is formed as an intermediate. Methyl groups can donate electron density to

stabilise the positive ion (this is known as an inductive effect).

E.g. Ethene reacting with hydrogen bromide

b. Polymerisation

A polymer is a very large molecule made up of long chains of smaller units (monomers) joined together. Artificial polymers include plastics and are

important materials for modern life.

Alkenes undergo addition polymerisation where the double bond breaks open to link together repeating units in a long chain.

E.g. polythene

Polypropene

Polychloroethene (PVC)

Polytetrafluoroethene

5. The properties of polymers

The properties of polymers depend on the chains of monomers from which they are formed. They can vary greatly.

Feature Property

The average length of the polymer chain

Branching of the chain

The presence of intermolecular forces between chains

Cross-links between chains

6. Polymer problems and solutions

(Refer to HSW box p136 – 137)

a. Problems

The production of synthetic polymers has saved many natural materials from destruction, but there is a down side to this.

i. Energy costs

Polymers are produced from fossil fuels. There are hidden energy costs to this. Around 4% of the worlds fossil fuel supply is used to generate the

electricity needed for polymer production.

ii. Resources used

There is a limited supply of fossil fuels with rising costs. This gives an incentive to produce polymers from different monomer units.

iii. Disposal problems

They aren’t easy to dispose of and cause a substantial waste problem. A variety of toxic gases are released when burnt, including hydrogen cyanide.

They also cause danger to wild animals.

iv. Carbon footprint The use of fossil fuels in polymerisation releases carbon thereby increasing

our carbon footprint. The quantity of non-biodegradable polymers is greatly increasing.

b. Solutions

i. Renewable energy sources

These could be used to generate the electricity required to make polymers.

ii. Reducing use of polymer products There is a move to reduce the amount of packaging used e.g. in food. Higher

quality plastics are being used in packaging to encourage reuse e.g. ‘bag for life’.

iii. Recycling

Many plastics are difficult to dispose of in a way that doesn’t damage the environment. Many synthetic polymers are non-biodegradable.

iv. Recycling thermoplastics

So far only thermoplastics (plastics that soften on heating) can be recycled.

If the plastic that is recovered after recycling is as good as the original, then effectively all the hidden energy from its original production has been saved apart from the energy used in recycling.

7. How are plastics recycled?

There are 2 main ways of recycling;

Type of recycling Key points

Mechanical recycling

Chemical recycling

8. Biodegradable polymers Biopolymers are now being produced through the modification of natural

polymers such as cellulose and starch. Bacterial fermentation is used to break down the polymer. The degradation may take 20-30 years but the

product will eventually be destroyed.


9. Energy recovery

Burning waste polymer products can be used to generate electricity, reducing

the need for fossil fuels. The incinerators used have special pollution controls.

Pyrolysis and gasification are both processes where polymer wastes are converted into energy-rich fuels.

Many of these processes aren’t very efficient at the moment.

10. Life cycle analysis


This is used to quantify the effect of polymers on the environment. An inventory is made of the materials and energy used as well as environmental

emissions. (Refer to table 1.7.5 p141)

Documents

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