Chemistry Module 1 Revision

Formulae, Equations and Amounts

Energetics and Enthalpy changes

Atomic Structure and the Periodic Table

Bonding

Organic Chemistry

The Alkanes

The Alkenes

Formulae, Equations and Amounts

Chemistry Basics Make up of atoms

IsotopesChemical Reactions

State SymbolsIons in solution

MolesRelative Atomic MassRelative Formula MassMoles and EquationsReactions with Gases

Empirical/Molecular FormulaMoles in Solution

ConcentrationYield/ Atom Economy

Make up of an Atom

Every atom is made up of Protons, Neutrons and Electrons.An element is a substance which cannot be broken into chemically simpler substances.In the periodic table an atom looks like this :

27 is the mass number which isthe total number of nucleons

13 is the atomic number, thenumber of protons

Isotopes

All atoms of an element have the same atomic number.

However some have different numbers of neutrons. These have no charge but affect the weight of the element.

Isotopes have the same chemical properties but different physical properties such as being radioactive.

For example Hydrogen has 3 isotopes,

All but the first one are very unstable and undergo radioactive decay

Chemistry Basics 1

Al27

13

Representing Chemical Reactions

Chemists use a simple method to show what is happening in chemical reactions. These must have the same number of each element on each side.

The total mass of the product is equal to the total mass of reactants.

State SymbolsEvery equation must include state symbols for each chemical.

Ionic EquationsSome reactions involve ions in solution. These are often shown as ionic equations where only the ions involved in the reaction are shown.The molecular equation would be:

The ionic Equation would be :

Bond BasicsThere are two main types of bond :

Chemistry Basics 2

2Mg + O2 → 2MgO

•Solid

s

•Liquid

l

•Gas

g

•Solution

aq

2NaOH ( a q ) + MgCl 2 ( a q ) → 2NaCl ( a q )

+Mg(OH) 2 ( s )

2OH -(aq) + Mg2+

(aq) → Mg(OH)2(s)

Ionic Bonds Covalent Bonds

Atoms lose or gain electrons so they have a full outer shell

Atoms share electrons so both have a full outer shell.

Metal – Non Metal Metal - Metal

Relative Atomic MassChemists use a relative scale to weigh atoms.The RAM is defined as the average mass of its isotopes compared with the mass of one atom of carbon-12.Lots of atoms are made up of several different isotopes. The average mass is the actual RAM, although this is usually very close to the major isotope.

Formula MassFor a compound find the sum of the RAM’s of all the atoms in the chemical formula. E.g.

Mr of CO2 = 12 + (16x2) = 44

In covalent compounds it is usually called the relative molecular mass.

The molar mass is the relative molecular or formula mass in grams per mole.

The Mole

The mole is used to compare ratios of atoms.A mole of any substance is the amount of substance that contains as many particles as there are atoms in 12g of carbon-12.One mole of any substance contains the Avogadro constant (6.02 x 1023) of particles.

How many atoms of zinc are there in 16.35g?

16.35 ÷ 65.4 = 0.25 moles0.25 x 6.02 x1023 = 1.505 x1023

Using moles you can work out the amounts of reactants needed for a reaction. E.g. to make one mole of MgO you need:

2Mg + O2 → 2MgO

2 mol + 1 mol → 2 mol2 x 24.3 + 32 → 2 x 40.3

Moles 1

Moles = Mass . Molar

mass

Reactions with Gases It is easier to measure the volume of gases than weigh them. This is easy becauseOne mole of any gas occupies 24dm3 at 298K and 1atm.

What volume of CO2 is produced by burning 6g of C?

C + O2 → CO2

1 mol + 1 mol → 1mol6 ÷ 12 = 0.5 moles0.5 x 24 = 12dm3

Empirical FormulaThis can be worked out using results from an experiment. E.g. Analysis of Aluminium Chloride shows it contains 5.8g of Al and 22.9g of Cl.

So the Empirical Formula is AlCl3.

However the molecular formula could also be Al2Cl6.

The only way to find the molecular formula is if you know the RAM.

Moles 2

Al Cl

Mass from analysis (g)

5.8 22.9

Molar Mass (mol-1) 27 35.5

Moles of atoms 0.215 0.645Ratio of moles 1 3

Moles in Solution When moles occur in a solution they are often in different concentrations. This is measured in

Moles per cubic decimetre (mol.dm-3)

How many grams of NaCl are needed to make 500cm3 of 0.150 solution?Mr (NaCl) = 23 + 35.5 = 58.5 g mol-1

0.15 x 58.5 = 8.8g (for 1dm3)4.4g (for 500cm3)

ConcentrationThe concentration, or molarity of a solution can be worked out using:

Other units of concentrations:Percentage by mass

Mass of solute / mass of solution

Percentage by volumeVolume of one component / total volume

Parts per millionThe formula for working out the number of moles in a solution is therefore:

Moles 3

Moles of Solute x 100volume of solution

Moles = vc .

1000

Yield of reactionsIn reality very few reactions actually produce the calculated amount of product. This is because:

The reactants may not be pure The reaction may not finish Product might be left on the apparatus Human error may have an affect

6.4g of copper oxide is reacted with excess sulphuric acid and 14.7g copper sulfate crystals are produced.Mr (CuO) = 63.5 + 16 = 79.5 g mol-1

Mr (CuSO4.5H20) = 63.5+32.1+4x16+5x18 = 249.6 g mol-1

So 79.5g of CuO should produce 249.6g crystals.Moles of copper oxide : 6.4/79.5 = 0.08molYield 100% would make 0.08x249.6 = 20.0gPercentage Yield = 14.7/20 x 100 = 73.5%

Atom Economy This is the idea that the important part of a reaction is not the overall yield but the amount of products when compared with the amount of reactants.

Therefore Atom Economy =

We use the relative formula masses of each atom.

It can be used to create more effective ways of creating drugs and chemicals with less waste and therefore providing a commercial

Mass of atoms in desired product x 100 Mass of atoms in reactants

Moles 4

Energetics and Enthalpy changes

Energy and EnergeticsMeasuring enthalpy changesHeat capacities and CalorimetersImportant Enthalpy changesUsing enthalpy changesBond enthalpiesUsing Bond Enthalpies

Energy ChangesEnergetics

This is the study of energy transfers between reacting chemicals and their surroundings. The study of the heat transfers during reactions is also called thermochemistry.

Exothermic and Endothermic ReactionsMany reactions need energy to happen, others give off energy. This is determined by what is happening to the chemical bonds.

Bond breaking requires energy while bond making releases energy

Exothermic Energy Changes

The energy released in the bond formation of the products is greater than the energy needed to break the bonds of the reactants.This is normally combustion reactions and neutralisations reactions.

Endothermic Energy Changes

The energy required to break the bond in the reactants is greater than that needed in the bond formation of the

products.

This is often seen in thermal decomposition reactions and photosynthesis.

SLIGHTLY EXOTHERMIC

EnthalpyAnother term for the heat of reaction is the enthalpy change of reaction.

The reaction in which these enthalpy changes happen is called the system. Everything else is referred to as the surroundings.

Some systems have physical boundaries between its surroundings, this is called a closed system. Sometimes this prevents energy leaving or entering the system, in which case it is isolated. However most remain open to energy transfers.

If we can measure the amount of energy leaving/ entering the system then using the idea that

Energy cannot be created of destroyedwe know what has happened in the reaction.

Amount of energy transferredChemists define the energy content of a system held at constant pressure as its enthalpy (H).This is impossible to measure but chemists can work out the change in enthalpy.

Enthalpy Changes

∆H = Hp r o d u c t s - H r e a c t a n t s

For endothermic reactions the enthalpy of reaction is positive, the energy content of the system has increased.

For exothermic reactions the enthalpy of reaction is negative, the energy content of the system decreases.

Enthalpy ChangesHeat Capacities

The heat capacity of an object is the amount of heat required to raise its temperature by 1K. Its unit is JK-1.

This however only applies to a single object.

The specific heat capacity is the amount of energy in Joules needed to raise the temperature of one kilogram of a particular substance by one Kelvin.

This also allows us to use the equation that tells us the energy transferred is equal to the mass of the object times the specific heat capacity times the temperature change.

Example:25cm3 of 1M HCl was mixed with 25cm3 of 1M NaOH in a coffee cup calorimeter. Both were at 22.5°C and the temperature rose to 29.3°C. Find the enthalpy change per mole of HCl assuming:

• It has the SHC of 4.2Jg-1K-1

• Their densities are 1gcm-3

Energy Transferred = mass x SHC x ∆TE = 50 x 4.2 x 6.7 = 1407J

Moles = vc = 1 x 25 = 0.025mol 1000 1000

∆ H = energy = 1407 = 56.3kJmol - 1

moles 0 .025

E = mc∆T

Enthalpy ChangesStandard Enthalpy changes

There are several different types of Enthalpy change which can be measured and calculated. However we need to keep the conditions constant for them all to remain relative. This means the reaction must be under standard conditions (298K and 1atm).It would be presented like this:

You MUST have the ± s ign

Things that affect theenthalpy change ofCombustion:•Number of bond to be broken

This depends on the size of the molecule•Type of bonds involved

Shows how the chemical make up affects a fuel.

±∆HѲ

Standard enthalpy change of formationThe enthalpy change for the formation of one mole of the compound from its constituent elements in their standard states under standard conditions.

Standard enthalpy change of combustionThe enthalpy change for the complete combustion of one mole of the substance under standard conditons

Standard enthalpy change of atomisationThe enthalpy change for the formation of one mole of gas atoms from the element under standard conditions.

Standard enthalpy change of neutralisationThe enthalpy change of reaction when one mole of acid is just neutralised by an alkali in their standard states at 25˚C in 1.0 Mol.dm-3 solutions

Enthalpy ChangesHarder Enthalpy Calculations

If we are told the standard enthalpy change of combustion of carbon is – 393.5kJmol - 1 , we can assume the enthalpy change for the decomposition of CO2 would be +393.5kJmol - 1 .

We can also draw a triangle of reactions in which carbon reacts with oxygen to form carbon monoxide. This reacts with more oxygen to form carbon dioxide.We know the ΔH of formation of CO2 is -393.5 kJ.mol-1

We know the ΔH of combustion of CO is -283 kJ.mol-1

Therefore -393.5-(-283) = -110.5 kJ.mol-1

Hess cyclesThis works on the idea that:The total enthalpy change for a reaction if independent of the route taken.

Remember each enthalpy figure is per mole, multiply it by the number of molecules in the reaction.

See page 44 for examplesCO(g) + ½O2 (g)C(Graphite) + O2

(g)

CO2 (g)

?-393.5 kJmol - 1 -2

83 kJmol-

1

Elements

Reactants

Products

ΔHfΔHf

ΔH

Enthalpy ChangesBond Enthalpies

This is useful to calculate the energy change in a reaction involving covalent bonds.

The energy needed to make/break a particular covalent bond is called the bond dissociation enthalpy.

Bond enthalpies are useful because:• They allow you to estimate the enthalpy changes in

reaction• They can be used to compare the strengths of

bonds between atoms• They help understand the structure of compounds

The calculated values vary slightly from the experimental values because the figures used are averages.

Mean Bond EnthalpiesThis is the average dissociation enthalpy of a particular bond over a large range of energies.

X = +1163.5 kJmol-1

E (C—H) = 1163.5 ÷ 4 = +415.9 k Jmol - 1

Atomic Structure and the Periodic Table

Measuring the Mass of an AtomMass Spectrometer Data

Electron arrangementFilling Orbitals and Shells

Chemical PropertiesPeriodic Table Blocks

Patterns in the Periodic Table

Measuring the mass of an atomMeasuring the mass of an Atom

Chemists want to be able to measure the Relative atomic/formula mass of elements and chemicals. They do this using a Mass Spectrometer, which works as opposite.

1. Vaporising the sample into a gaseous state.

2. Bombarding the vapour with electrons, forming positive ions.

3. Accelerating the ions using an electric field.

4. Passing the ions through a velocity selector ensuring they all travel at equal speeds

5. Passing the ions through a uniform magnetic field. The heavier ions are deflected less than the lighter ions. The magnetic Field strength is slowly increased.

6. The detector counts the ions passing through the machine and can show how many ions of each mass: charge ratio there are.

Uses of Mass SpectrumsMeasuring the mass of an Atom

We then obtain a mass spectrum which shows the mass of ions detected and the relative abundance.

We can see there are two regions, one for chlorine atoms and one for Cl2 molecules. To find the RAM of chlorine find the average of the Ions detected, giving us 35.5 in this case.

Mass Spectrometry can also be used for:

• Finding the age of materials using carbon dating , by analysing the amount of Carbon-14 against the Carbon-12.

• Detecting drug cheats in sports, analysing the amount of testosterone against natural epitestosterone.

• In space they can be used to work out the atomic make up of gases in the atmosphere.

• Identifying molecules with potential use as drugs for pharmaceuticals

35 37 70 72 740

20

40

60

80 Mass Spectrum of Chlorine

Mass of Detected Ions

% a

bu

nd

an

ce

The Electronic Structure of AtomsElectron Emission Spectrum

When a gas is heated and electrically charged it gives off a light. When passed through a spectrum it gives a pattern which is specific to each element.The light energy can be translated into the ionisation energies needed to remove electrons from an atom.By looking at ionisation energy data we can work out the electron structures in atoms.

Ionisation Energy

The Ionisation Energy is the:Energy required to remove one mole of

electrons from one mole of atoms in the gaseous state.

This is always endothermic. There are more than one ionisation energy, every electron removed requires more energy.

Evidence for Electron ShellsBy looking at the series of successive ionisation energies of Sodium:

First ionisation energy is lowest, removed from outer shell.Large increase between 1st and 2nd I.E. And 9th and 10th I.E as a shell is removed.Electrons nearer nucleus attracted more strongly.You can use this to work out what group the element is in by number of I.E’s to shell change.

0 2 4 6 8 102

3

4

5

6

Lo

g I.

E.

The Electronic Structure of AtomsEvidence for Electron Sub-Shells

By looking at the first ionisation energy of successive elements:

Generally ionisation energies increase across a period.The slight decrease between atomic numbers 4 and 5 and 7 and 8.This suggests there are several sub-shells which have different energy levels.

Scientists have been able to look at the patterns and work out that there are four sub shells.Electrons are also paired into Orbitals which are made up of electrons with opposite ‘spin’.

Each orbital has a specific shell shape.

The shells fill up in order of the lowest energylevels, the shell is notalways complete whenthe next electron shellis started.

0 2 4 6 8 10 12 14 16 18 200

500

1000

1500

2000

2500

1s

t I.E

Sub Shell Electrons Orbital PatternS 2 ↑↓P 6 ↑↓ ↑↓ ↑↓D 10 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓F 14 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

6s 6p 6d

7s

Electrons and OrbitalsElectron Density Maps

Each electron type has a different shape for the space it is most likely to be in. They can be described as electron clouds. By drawing electron density maps we can start to understand the size and shape of atoms. S Orbitals are spherical and layer themselves round the nucleus.P Orbitals are dumb-bell shaped and sit on three axis, at right angles from each other.The picture below shows an atom with electronic configuration 1s2 2s2 2p6.

Electronic ConfigurationsWe can easily describe the electronic structure of atoms using simple notation. Just follow along the shell filling pattern and write how many in each sub shell. Remember to write in SHELL order:

The Periodic Table is divided into blocks which show the pattern.

Bromine →1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5

Patterns in the Periodic TablePeriodicity means the patterns seen in the properties of elements in the periodic table.

Atomic RadiusACROSS A PERIOD THE RADIUS DECREASES:More Protons in nucleus means the outer shell electrons are pulled in more.DOWN A GROUP THE RADIUS INCREASES:Each New shell means the proton increase is offset by the inter-shell repulsion

Ionisation EnergyTHE 1ST I.E. INCREASE ACROSS A PERIOD:The proton number in the nucleus means there is a stronger electrostatic attraction on the outer shell.THE 1ST I.E DECREASES DOWN A GROUP:An extra shell is added placing the electrons further from the nucleusExtra shells means the outer shells are repelled

more.

Melting PointsThe melting point temperature of an element depends on its structure and bonding.Melting temperatures rise from group 1-3 as the number of shared electrons in the ‘sea of electrons’ holding the metal together gets larger.Group 4 metals form giant covalent structures with the highest melting points.Groups 5-0 form simple molecules with low intermolecular forces, these have low melting points.

2 4 6 8 10 12 14 16 180

1000

2000

3000

4000

Me

ltin

g T

em

p /°

C

Patterns in the Periodic TableExceptions

Some elements break the trend.

1st Ionisation Energies of Beryllium greater than Boron: (Also Magnesium and Aluminium)

Beryllium’s outer electron is in the 2s sub-shell Boron’s outer electron is in the 2p sub-shell.This means boron’s outermost electron is easier to remove because the 2p sub-shell has a higher energy level than the 2s sub-shell.

1st Ionisation Energies of Nitrogen greater than Oxygen: (Also Phosphorus and Silicon)

Both have their outer electrons in the same 2p sub-shell. (same energy level)Nitrogen’s outer shell has 3 unpaired electrons in it’s 3 orbitals.Oxygen has 2 unpaired and one paired electron in it’s 3 orbitals.These paired electrons repel each other, making it easier to remove an electron from oxygen.

BONDING

Ionic Bonding

Giant ionic Lattices

Evidence for ions

Lattice Energy

Covalent Compounds

Metallic Bonding

Ionic BondingChemical Bonds

A chemical bond is the force which holds together two bonds. Atoms can lose, share or gain electron when they react.

Evidence for ions

Physical properties of ionic compounds show strong forces of attraction between ions. They also conduct electricity when in solution or molten

Electron density maps show zero election density between ions so there must be complete electron transfer

Electrolysis proves ions exist because of positive ions are attracted to the cathode and negative ions to the anode.

Forming Ionic BondsThey are normally formed between metals and non-metals. Oppositely charged ions are attracted to each other by an electrostatic force called an ionic bond.The ions have full outer shells so you can work out the charges on the ions from their electronic structures.

The transfer of electrons from one atom to another is the ionic (or electrovalent) bonding and can be represented using dot and cross diagrams.

Giant Ionic LatticesTrends in Ionic Radii

The ionic radius is the radius of an ion in a crystal. The radius of a positive ion is smaller because the remaining electrons are more strongly attracted to the nucleus. Negative ions are also bigger than their atoms.

Lattice Structures

Ionic Compounds form giant ionic lattices which are crystals. The lattice structure is the arrangement of ions that has the strongest attractive forces between oppositely charged ions.

If X-rays are passed through a crystal they are scattered or diffracted by the electrons in the ions in the structure. The pattern is known as an electron density map.

Types of Lattice

There are lots of different types of lattices. The coordination number is the number of atoms ‘touching’ each other atom.

In the case of Sodium chloride this is 6.

This can also be used to provethe formula:1 Sodium Ion is next to 6 chlorineions. However each of these isshared with 6 other ions soeffectively it joins with 6 x 1/6 Cl- ions.

So the formula is NaCl

Lattice EnergyChemical Bonds

The formation of an ionic crystal is exothermic.The lattice energy is

The energy released when one mole of an ionic crystal is formed from its ions in the gaseous state

under standard conditionsThis process can be broken down into several stages:

•Atomisation of the Metal•Atomisation of the Non- Metal•Ionisation of the Metal•Electron Affinity of the Non-Metal•Formation of the Crystal from Gaseous State

These all equal the heat of formation of the compound.

All the enthalpies but one will be given, so this can be worked out as we know:

Clockwise = Anticlockwise

Finding Lattice Energy of CaODraw a Born- Haber cycle (like a Hess cycle):

Ca O

Ca2+O2-

Ca+O-

CaO

Ca2O2

∆Hf(CaO)

∆Hat(O2) ∆Hat(Ca2)

∆HEA(O)

∆HEA(O-)

∆Hi1(Ca)

∆Hi2(Ca+)

∆HLat(Ca2+O2-)

StabilityPatterns in Lattice Energies

Charge Density is a measure of how concentrated the charge is. A high CD is found in smaller ions with higher charges.

In smaller ions the attractive force between the electrons and the nucleus is greater because they are closer.

The size of the force of attractionof one ion on another can becalculated using Coulomb’s Law,When q1 and q2 are the ionic charges and r is the distance between them.

StabilityThe Born-Haber cycle can be used to calculate the standard enthalpy of formation of compounds which do not normally exist. For example we can work out that MgCl would have ∆Hf of -94kJmol-1. However MgCl2 has ∆Hf of -641kJmol-1 . This means MgCl 2 is most stable.

Different Lattice EnergiesChemists calculate Lattice energies using measured enthalpies and the Born-Haber cycle.They can also calculate it using Coulomb’s Law.

The results are similar for the Sodium halides, however the silver halides are more exothermic than theory would suggest.Coulombs law relies on the ions being separate and spherical. Experiment suggests some electrons are shared in the silver, whereas sodium shows pure ionic bonding. This is supported by their melting points.

F = k q1q2

r2

CompoundLattice Energy /kJmol-1

Born – Haber Theoretical

NaF -918 -912

NaCl -780 -770

NaBr -742 -735

NaI -705 -687

AgF -958 -920

AgCl -905 -883

AgBr -891 -816

AgI -889 -778

PolarizationPolarization of an ion is the distortion of its electron cloud away from completely spherical.

The nucleus of a cation (+) will attract the outer electrons of an anion (-). This means the cation has polarizing power and the anion is polarizable.

The polarizing power of a cation depends on it’s charge density:

• A small cation is more polarizing than a large one, the small of ionic radius means it

has a stronger attraction.•A cation with a large charge is more

polarizing than one with small charge because they have greater attraction.

The polarizability of an anion depends on its size alone:•A large anion is easily polarized as its outer

electrons are further away.

Figure A is a wholly ionic bond

Figure B is the beginning of polarization

Figure C is extreme polarization where it starts to represent a covalent bond.

Covalent BondingForming Covalent Bonds

A covalent bond is formed when two electrons are shared between two atoms, their electron clouds overlap and they both effectively have a full outer shell.

The overlapped region of high electron density attracts both nucleus’s keeping the atoms together.

They are very strong electrostatic attraction.The distance between the nuclei is the bond length.

Dative covalent bonds are formed when both shared electrons come from one atom.

Atoms can share more than one pair of electrons, a double bond or triple bond are formed when multiple pairs are shared.

A lone pair is two outer shell electrons that are not shared, they can affect the shapes of molecules.

Dot-and-cross DiagramsThey can be drawn as dot and cross diagrams:

Covalent BondsGiant Atomic StructuresThese from distinctive atomic crystals. The whole lattice can be thought of as a giant molecule, such as diamond.They are very hard and have high melting points showing their high melting points.

Electron Density MapsFor a simple covalent compound the bonds are highly directional. They show that there is high electron density between the covalently bonded atoms.

H HOx x

x x

x x

oo H HO

Metallic BondingMetals consist of giant lattices of metal ions in a sea of delocalised electrons; The ions vibrate about fixed point and are held in place by the surrounding electrons.

The outer electrons of the metal become delocalised, and are no longer associated with that particular atom.

Metallic Bonding is the strong attraction between the ions and the electrons.

Typical Characteristics of MetalsElectrical Conductivity

The delocalised electrons are free to move when a current is passed through it, this movement of charged particles completes the circuit.

Thermal conductivityThe electrons transmit kinetic energy through the metal by collision with each other.

High Melting TemperaturesThe positive ions are strongly held together by the attraction of the electrons, this means it takes lots of energy to break the bonds and allow the particles to move in a liquid state.

Malleability and DuctilityMetals can be hammered into shape (malleable) or stretched into wires (ductile) because the ions are in layers which can be forced to slide across each other whilst still surrounded by electrons.

Organic Chemistry

Risk and HazardHomologous SeriesNaming and Drawing

Introductory Organic ChemistryTypes of Chemistry

Organic ChemistryThis is the study of carbon compounds. There are over 7 million compounds formed with carbon, scientists study their properties and uses.

Inorganic ChemistryThe study of the 91 naturally occurring elements and their compounds, includes a few simple carbon structures like CO2.

Hazard and RiskMany organic substances are potentially harmful. Two definitions you need to know are:Hazard: The potential of a substance to do harm. These

are represented using hazard symbols

Risk: The chance that a substance will do harm. This can be assessed and reduced if the risk is unacceptable.

How to reduce risks:•Using less material•Using lower concentrations•Using electric heaters not naked flames•Using a fume cupboard•Using similar but less hazardous materials.

Organic compoundsCarbon forms four bonds

Carbon always forms four covalent bonds so as to fill its outer shell.It is unique because it can form covalent bonds with other carbon atoms and other atoms at the same time.Carbon can form many different shapes and patterns are often seen.

Classifying Organic CompoundsThe first way to describe an organic compound is on the arrangement of it’s carbon chainAliphatic molecules contain straight or branched skeletons.

Alicyclic molecules are closed rings of carbon atoms which contain multiple carbon bonds.Arenes are all derived from benzene and are hexagonal carbon rings. (A2 Course, I don’t think so...)

Organic Compounds also fall into a carbon family, or homologous series. Their physical properties are similar but melting temperatures and viscosities increase with more carbon atoms. Each is characterised by a specific functional group.

Functional GroupsH. Series Formula Example

Alkanes CnH2n+2 C2H6

Alkenes CnH2n C2H4

Alkynes CnH2n-2 C2H2

Alcohols CnH2n+1OH C2H5OH

Aldehydes CnH2nO CH3 CHO

Ketones CnH2nO CH3 COCH3

Carboxylic Acids CnH2nO2 CH3COOH

C H

OH

CC

CC

CO

H

C

O

CO

OH

Naming and DrawingIUPAC System

There is an international system for naming compounds.

The prefix’s in the table are used depending on the number of carbon atoms.

The suffix depends on the homologous group.Numbers are used to represent which carbon

is affected.

There are several ways a chemical can be represented:Empirical Ratio Between elementsMolecular How many atoms in a moleculeStructural Shows atoms and functional GroupsDisplayed Shows bonds as sticks

3-ethyl-4, 4-di-methylhex-2-ene

Structural: CH3CH2CH(CH2CH3)C(CH3)(CH3)CH2CH3

Molecular : C10H22

Prefix CMeth- 1Eth- 2Prop- 3But- 4Pent- 5Hex 6Hept- 7Oct- 8

C

C

C

C

C

C

C C

C

H

H

HH

HH

HH

HH

HH

H

H

H

HH

H

H

CHH

H

THE ALKANESHydrocarbons

Using Crude OilChemical PropertiesChemical Reactions

General PropertiesAlkanes are a family of saturated hydrocarbonsAll Hydrocarbons they are:

•Insoluble in Water•Burn in oxygen

They generally get more dense and boil at higher temperatures the more carbons in the chain.

Alkanes with the same formula but different structural formulas are structural isomers. E.g. butane has formula C4H10 but there are 2 different ways to make this chemical.

CRUDE OILAlkanes come from crude oil. This is then broken up

using fractional distillation. Different amounts of each fraction are needed world

wide mostly lighter fractions.

Scientists have worked out how to break the longer chains into shorted easier to burn ones. This is called cracking.

Reformation is similar to this. It uses a catalyst such as platinum to break down long chains into shorter ones and arenes (rings for the chemical industry).

ALTERNATIVE FUELSCrude oil is not a sustainable resource.It’s use is also thought to be causing global warming.

For these reasons we need new reliable and efficient ways to get energy. Some ideas include:

•Bio fuels made from quick growing crops•Burning Hydrogen gas•Combining Hydrogen and Oxygen in fuel cells.

However this is not yet efficient enough to produce more electricity than the whole process needs.

Chemical PropertiesApart from combustion the alkanes are a very unreactive family. They are:

•Non-corrosive•Harmless to the skin•Un-affected by conc. acids at room temperature•Not affected by oxidising agents•Do not react with metals

Nearly all reactions the of the Alkanes are due to the formation of free radicals. #

CRACKINGWhen Alkanes are heated to high temperatures they split into smaller molecules. Cracking is thermal decomposition which involves the breaking of C – C bonds. It always gives alkenes with C = C bonds.

BREAKING BONDSWhen covalent bonds are broken there are two ways the electrons can be shared out. Homolytic FissionThis involves the equal sharing of the electrons. We use curly arrows to show what is happening.This results in two uncharged atoms, called free radicals. They are extremely reactive.

Heterolytic FissionThis involves the unequal sharing of the electrons. This results in two charged particles. This normally happens when the bond already has a degree of polarity

Cl ― Cl → Cl• + Cl •

H ― Cl → H+ + Cl-

Chemical ReactionsCOMBUSTION

Alkanes burn in excess oxygen to produce carbon dioxide and water. This is complete combustion.However when burned in a limited supply of oxygen Alkanes form carbon monoxide under incomplete combustion.

Burning hydrocarbons releases greenhouse gases, this is one reason why scientists are looking for new ways to produce energy.

SUBSTITUTIONThis is a replacement reaction. For example a hydrogen atom can be replaced by a different atom or group.

This is called free radical substitution. A radical is a species with an unpaired electron.Alkanes undergo substitution by halogens at about 300°C or in ultraviolet light.

These three stages form a reaction mechanism:Initiation UV light breaks the covalent bond by homolytic fission to produce free radicals.

Propagation Where the substitution products are made in two consecutive steps. This also regenerates the free radical that caused the original reaction and can take place multiple times.

Termination Can only occur when two radicals meet and mutually annihilate each other. The radical are removed and the reaction stops.CH4 + Cl2 → CH3Cl + HCl

Cl ― Cl → Cl• + Cl•

Cl• + CH4 → •CH3+ HCl•CH3 + Cl2 → CH3Cl + Cl•

Cl• + Cl• → Cl2•CH3 + •CH3 → C2H6•CH3 + Cl• → CH3Cl

THE ALKENESDouble bondsGeometric IsomerismChemical ReactionsHow do organic reactions happen?Polymers

Double BondsAlkenes are another family of hydrocarbons, however they are unsaturated.They are mostly obtained by the process of cracking and are very important to the chemical industry.The main difference is way the carbon atoms are bonded:In an Alkane a single sigma bond holds the atoms together as a covalent bond. This is symmetrical to the central axis of the molecule and both ends are free to rotate.In an Alkene there are two bonds, the sigma bond and a pi bond. This has concentrated electron density above and below the axis meaning it does not allow rotation.

Geometric IsomersBecause they cannot rotate different groups can occur on each side of the molecule, giving it different chemical and physical properties.

Cis-trans isomerismThe old system based on the location of the functional groups.•If the functional groups are on the same side they are CIS.

•If they are on opposite sides they are TRANS.

E-Z isomerismThis system works for all isomers. Each group is based on atomic number.

•High ranks are on the same side = Z•High ranks are on opposite sides = E

General properties

EZ

Addition ReactionsThe C=C bond makes alkenes much more reactive when chemicals are added. The main ones to know are:

Addition of HydrogenThis is called Catalytic hydrogenation. It is used to turn unsaturated vegetable oils in to saturated fat (e.g. making margarine). It needs to be done at about 200C with a nickel catalyst.

Reaction with Acidified Potassium Manganate (VII)When diluted in dilute sulfuric acid and reacted with a alkene it results in oxidation the alkene to form a diol. (2 -OH groups) . The potassium manganate (VII) changes from purple to colourless.This is not the industrial way to produce diol’s but is a good test for alkenes as alkanes would not react.

Addition of HalogensThis produces a disubstituted halogenoalkane. The reactivity of the halogens decreases down the group.They can occur at room temperature, suggesting free radicals are not involved.When an alkenes is mixes with bromine water the solution becomes colourless because the product is 2-bromoethanol and 1,2-dibromoethane.

THE MECHANISMElectrophilic Addition across the C=C bond has a two-steps:

•Electrophilic attack on the double bond, one carbon atom gains a positive charge forming an ion called a carbocation.

•Nucleophilic attack on the carbocation where the remaining bromine ion is bonded.

Alkene Reactions

CH4 + H2 → CH6

Addition of Halogen HalidesThis produces a monosubstituted halogenoalkane, the reaction occurs at room temperature.The mechanism is the same as for the addition of halides, however in longer alkanes chains there are multiple possible products.

The major product can be determined using Markovnikov’s rule:When HX is added across an asymmetric double bond, the major product formed is the molecule in which hydrogen adds to the carbon atom with the most hydrogen atoms

already attached to it.

Evidence for Reaction MechanismsIf these addition reactions are done at the same time in the presence of competing nucleophiles then a mixture of two products are formed.

The two products can only be formed it the first step results in the formation of a carbocation, to which a negative ion can bond.

Br2 alone

Br2 with Cl- ions

andCl

Alkene Reactions

Addition polymerisationA polymer is formed when a very large number of monomers (small molecules) join together to form a chain.Alkenes undergo additional polymerisation to form a polyalkane.The C=C bonded monomer is a repeat unit, this is replaced by a single bond and side links. The formulae uses a letter n to show how many links there are.

Chloroethene Polycholorethene

A monomer alkene can be identified from the polymer, it includes a C=C double bond and four groups attached to two carbon atoms.

Properties of PolymersThe properties depend on several things:

•Average length of chain•Inter molecular forces between the chains•Cross-links between the chains.•Branching on the chain

Polymer lifecyclePolymers are made from chemicals from crude oil so

have high energy production costs and use non-renewable resources.

They are also hard to dispose of, when burnt they give off toxic gases and do not biodegrade for hundreds of years.

We therefore recycle them, this still has some energy costs but is much better overall.

Some areas use energy recovery, where they are burnt at high temperatures, no gases are given off and electricity can be produces from the heat.

Polymerisation

nn