AS Level Chemistry | Unit 1 Revision Guide

1 Joel Biffin

AS Chemistry

Unit 1: Foundation Chemistry

Joel Biffin

2 Joel Biffin

Atomic Structure [1.1]

Amount of Substance [1.2]

Bonding [1.3]

Periodicity [1.4]

Organic Chemistry [1.5]

Alkanes [1.6]

Unit 1: Foundation Chemistry Atomic Structure [1.1]

3 Joel Biffin

Unit 1 | Atomic Structure [1.1] Fundamental Particles

Particle Name Relative Charge Relative Mass

Proton +1 1 Neutron 0 1

Electron -1 5.45x104

Mass Spectrometry

- Used to determine relative abundance of a species, relative isotopic mass & relative molecular mass (Mr).

- The following takes place in a mass spectrometer:

> Ionisation Electron gun fires electrons at gaseous atoms, removing an electron thus creating gaseous ion

> Acceleration Ions are accelerated using an electric field and are focused by a small slit (exit point)

> Deflection A strong electromagnet creates a magnetic field, deflecting ions. Deflection depends on the m/z ratio (Mass/Charge) the greater the m/z the more deflection. Only ions with a specific m/z will pass through the spectrometer.

> Detection Ions hit a metal plate, pick up an electron thus inducing a current. The larger the current, the greater the abundance of species.

Electron Arrangement

- s sub-shell contains a max. of 2 electrons (i.e. 1s2)

- p sub-shell contains a max. of 6 electrons (i.e. 2p6)

- d sub-shell contains a max. of 10 electrons (i.e. 3d10)

M(g) -> M+(g) + e-

Unit 1: Foundation Chemistry Atomic Structure [1.1]

4 Joel Biffin

- Electrons in same sub-shell prefer to occupy different orbitals if possible (e.g. in 2p3 1 e- in each orbital.

- When forming ions, electrons are lost from 4s before 3d i.e. Fe2+ [Ar] 3d6

N.B. Chromium (Cr) and Copper (Cu) are exceptions to the expected patterns

Chromium 1s2 2s2 2p6 3s2 3p6 4s1 3d5 Copper 1s2 2s2 2p6 3s2 3p6 4s1 3d10

Definitions

Mass Number The number of neutrons and protons (nucleons) in the nucleus of an atom

Isotopes Atoms with the same atomic (proton) number but a different number of neutrons

1st Ionisation Energy The enthalpy required to remove one mole of electrons from one mole of gaseous atoms to produce 1 mole of gaseous 1+ ions

2nd Ionisation Energy The enthalpy required to remove one mole of electrons from one mole of gaseous atoms to produce 1 mole of gaseous 2+ ions

Unit 1: Foundation Chemistry Amount of Substance [1.2]

5 Joel Biffin

Unit 1 | Amount of Substance [1.2] Mole Calculations

- Working with known species or solids: e.g. Work out the number of moles in 54.4g of Sodium Chloride powder. Moles = 54.4 / 58.5 = 0.930mol;

- Working with aqueous solutions: e.g. Work out the concentration of aqueous Sodium Chloride when 0.930mol of NaCl powder is added to 200cm3 of water. Concentration = 0.930 / (200x10-3) = 4.65mol.dm-3;

- Working with gaseous substances: e.g. How many moles of Potassium Chlorate (V) must be heated to give 1.00dm3 of O2 at 293K and 10Pa n = (P.V) / (R.T) = (10 x 1.00) / (8.31 x 293) = 4.11x10-3mol;

Yield & Atom Economy

- Percentage Yield is a measure of the efficiency of an experiment.

- Percentage Atom Economy is a measure wasted atoms used to create the product.

Formulations

- Working out the Empirical Formula when given masses of atoms in a compound.

X Y Mass

Moles (Mass/Mr) Ratio (Moles/SmallestMoles)

Atoms in Formula

Unit 1: Foundation Chemistry Amount of Substance [1.2]

6 Joel Biffin

- Working out Molecular Formula when given Empirical Formula and Mr.

Definitions

Relative Atomic Mass The mass of an atom of an element compared to 1/12th the mass of an atom of Carbon-12

Relative Molecular Mass

The mass of a molecule compared to 1/12th the mass of an atom of Carbon-12

Avogadro Constant The number of Carbon atoms in 12g of Carbon-12

Empirical Formula Formula showing the simplest ratio of atoms of each element in a molecule

Molecular Formula Formula showing the actual number of atoms of each element in a molecule

Unit 1: Foundation Chemistry Bonding [1.3]

7 Joel Biffin

Unit 1 | Bonding [1.3] Bonding & Structure

- Ionic

> Bonding Giant Ionic Lattice of +ve and -ve ions with strong electrostatic forces of attraction between oppositely charged ions.

> Melting & Boiling Point High melting point due to strong electrostatic forces of attraction between oppositely charged ions which require large amounts of heat energy to overcome.

> Electrical Conductivity Good conductor of electricity when molten or in solution as ions are free to move and do so in the same direction when a voltage is applied. Poor conductor of electricity when solid as ions are held in a fixed lattice and furthermore are not free to move.

- Simple Covalent

> Bonding A pair of shared electrons between two atoms.

> Melting & Boiling Point Melting and boiling points generally low, but they depend on the strength of the intermolecular forces of attraction (van der Waals < permanent dipole-dipole attraction < hydrogen bonding).

> Electrical Conductivity Simple covalent molecules are poor conductors of electricity as electrons are not free to move.


8 Joel Biffin

- Dative Covalent (Co-ordinate)

> Bonding One atom donates both electrons to create a shared pair of electrons between two atoms.

Other properties are just like simple covalent.

- Macromolecular Covalent (e.g Diamond, Graphite, Silicon Dioxide) > Bonding

Very many, very strong covalent bonds (shared pairs of electrons).

> Melting & Boiling Point Very high melting & boiling points due to very many, very strong covalent bonds which require large amounts of heat energy to break.

> Electrical Conductivity Giant molecules are poor conductors of electricity as electrons are not free to move. Graphite is the exception to this rule as there is one delocalised electron for each carbon atom which is free to move between layers and electrons do so in the same direction when a voltage is applied.

- Metallic

> Bonding +ve metal ions are held in a fixed lattice surrounded by a sea of delocalised electrons.

> Melting & Boiling Point Melting and boiling points are usually high as there are strong electrostatic forces of attraction between +ve metal ions and the sea of delocalised electrons.

> Electrical Conductivity Metals are good conductors of electricity as the sea of delocalised electrons are free to move and do so in the same direction when a voltage is applied.

Polar Bonds

- Electronegativity is the power of an atom to attract the bonding pair of electrons in a covalent bond.


9 Joel Biffin

- Polar bonds are created by a more electronegative species forming a covalent bond with a less electronegative species (two species with a large difference in electronegativity).

> e.g. HCl H+ - Cl-

- Trends in electronegativity are as follows:

> Decrease as you move down the periodic table because the bonding pair of electrons are further

away from the nucleus and there is more shielding by full energy levels.

> Increase as you move across the periodic table because there is an increase in nuclear charge meaning that the bonding electrons are more strongly attracted to the nucleus of the atom.

N.B. Noble gases are not electronegative as they do not form covalent bonds.

Intermolecular Forces of Attraction

- Van der Waals Forces

> Temporary dipoles are created in all molecules (polar or non-polar) due to an uneven electron distribution caused by the movement of electrons.

> The larger the molecule, the more electrons it contains which furthermore increases the strength of the van der Waals forces between molecules.

> For similar sized molecules, shape can affect the strength of van der Waals - the more linear a molecule the greater its surface area and furthermore van der Waals forces are stronger.

- Permanent Dipole-Dipole Attraction

> Polar molecules (i.e. HCl but NOT CCl4) have an overall dipole due to the sum of the polar bonds.

> Between polar molecules there is attraction between the + of one molecule and the - of another molecule.

- Hydrogen Bonding

> Hydrogen bonding is a form of permanent dipole-dipole attraction between an H atom and an electronegative (N, O or F) atom.

> The lone pair from the N, O or F atom is attracted to the H nucleus of another molecule.

> Contrary to its name, hydrogen bonding is an intermolecular force not a bond.


10 Joel Biffin

Definitions

Ionic Bonding Strong electrostatic attraction between oppositely charged ions

Covalent Bond Shared pair of electrons between two atoms

Dative Covalent Bond Covalent bond in which shared pair of electrons comes from just one atom (or -ve) ion

Metallic Bond Electrostatic attraction between a lattice of +ve ions and delocalised electrons

Electronegativity Power of an atom to attract the bonding pair of electrons in a covalent bond

Polar Bond Covalent bond in which electron distribution I not symmetrical giving + and - ends

Polar Molecule Molecule in which the electron distribution is not symmetrical giving + and - ends (CCl4 is NOT a polar molecule)

Hydrogen Bonding Strongest intermolecular force of attraction wich arises between a hydrogen atom bonded to one of the most electronegative atoms (F, O or N) and the lone pair of another (F, O or N)

Unit 1: Foundation Chemistry Periodicity [1.4]

11 Joel Biffin

Unit 1 | Periodicity [1.4] Atomic Radius of Period 3

- Decrease as you move across the period due to:

> Electrons are in the same principle energy level.

> Proton Number/Nuclear Charge increases across the period.

Therefore, electrons are more strongly attracted to nucleus.

1st Ionisation Energy

- Generally there is an increase in 1st Ionisation Energy due to:

> Proton Number/Nuclear Charge increases across the period, and furthermore atomic radius decreases.

Therefore, more energy is required to remove an electron as you move across the period.

- There are two exceptions to this general trend:

> 1st IE of Al < Mg due to:

~ Als outer electron is in a different sub-shell (3p). ~ 3p electron is further from the nucleus than the 3s electrons is. ~ 3p sub-shell has greater shielding by full shells of electrons.

Therefore, less attraction between nucleus and 3p electron thus less energy is required to remove electron.

> 1st IE of S < P due to:

~ Ps 3p sub shell has 4 electrons (3p4) meaning there is 1 orbital with a pair of electrons.

Therefore, electron-electron repulsion means that less energy is required to remove electron.

0

200

400

600

800

1000

1200

1400

1600

Na Mg Al Si P S Cl Ar

1st

Ioni

satio

n En

ergy

/ k

Jmol

-1

0

0.02

0.04

0.06

0.08

0.1

0.12

0.14

0.16

0.18


Ato

mic

Rad

ius

/ nm

Unit 1: Foundation Chemistry Periodicity [1.4]

12 Joel Biffin

Melting & Boiling Points

- Metals (Na, Mg, Al): Increase as you move across the period due to:

> Electrostatic attraction between +ve metal ions and delocalised electrons.

> More delocalised electrons as you move across - more attraction between ions and electrons.

Therefore, stronger attraction between +ve ions and electrons requires more energy to overcome.

- Macromolecular (Si): High due to:

> Very many, very strong covalent bonds.

Therefore, more energy is required to break bonds.

- Simple Covalent (P4, S8, Cl2) & Monotomic (Ar): generally decrease due to:

> Strength of van der Waals forces between molecules are effected by the number of electrons in a molecule and the size of the molecules.

> S8 does not follow the general trend because there are more electrons in S8 than in P4.

Therefore, energy required to overcome van der Waals forces between molecules depends on size and number of electrons of a molecule.

Electrical Conductivity

- Increase in electrical conductivity across metals of period 3 (only metals conduct) due to:

> Delocalised electrons are free to move and do so in the same direction when a voltage is applied.

> As you move across the period there is one more delocalised electron per atom of the metal.

> Ions that are more +vely charge are smaller and furthermore provide less resistance when there is a current.

0

0.05

0.1

0.15

0.2

0.25

0.3

0.35

0.4


Elec

tric

al C

ondu

ctiv

ity

0

500

1000

1500

2000

2500

3000

3500

4000

4500


Mel

ting

& B

oilin

g Po

ints

/ K

Unit 1: Foundation Chemistry Alkanes [1.6]

13 Joel Biffin

Unit 1 | Organic Chemistry [1.5] Organic Nomenclature

Homologous Series Functional Group Prefix/Suffix

Alkanes

-ane

Alkenes

-ene

Haloalkanes

Chloro- Bromo- Iodo-

- Molecules can sometimes have attached molecular groups - these are called Alkyl groups.

> Methyl- CH3 > Ethyl- CH3CH2 > Propyl- CH3CH2CH2

Structural Isomerism

- Chain Isomerism

> Molecules where the carbon chain is different i.e. with an alkyl group

- Position Isomerism

> Molecules where functional groups are found in different positions i.e. -1-, -2-, -3- etc.

- Functional Group Isomerism

> Compounds that have the same molecular formula but contain different functional groups i.e. Aldehydes & Ketones Carboxylic Acids & Esters Alkenes & Cycloalkanes


14 Joel Biffin

Definitions

Homologous Series

Series of compounds which have: 1- same general formula 2- same functional group 3- a difference of CH2 between members 4- gradual change in physical properties

Functional Group Atom or group of atoms which, when present in different molecules, causes them to have similar chemical properties

Structural Isomerism Existence of two or more different molecules with the same molecular formula but different structural formulae

Hydrocarbon Molecule made up of carbon & hydrogen atoms ONLY

Saturated Molecule has only single bonds

Structural Formula Shows the unique arrangement of atoms in a molecule in a simplified form without showing all the bonds

Displayed Formula Shows ALL the bonds present in a molecule


15 Joel Biffin

Unit 1 | Alkanes [1.6] Fractional Distillation of Crude Oil

- Fractional Distillation is used to separate the hydrocarbons of different chain lengths.

- Due to van der Waals forces between hydrocarbons:

> Longest chain lengths are tapped off at bottom.

> Highest chain lengths are tapped off at top.

> N.B. branched chain alkanes have lower boiling points than their corresponding straight alkane.

Cracking

- Shorter chained hydrocarbons are more in demand so cracking is used to break down less useful hydrocarbons into more useful, shorter chained hydrocarbons.

- Generally: Long Chained Alkanes -> Shorter Chained Alkanes + Alkenes + H2

- Shorter alkanes have use as fuels and petrochemicals.

- Alkenes are used in the manufacture of polymers.

- Thermal Cracking

> High temperatures (900oC) and pressures (70atm).

> The higher the temperature, the lower Mr alkenes.

> The reaction has a free radical mechanism.

- Catalytic Cracking

> Cracking can occur in the presence of a catalyst at 450oC.

> The catalysts are zeolites.

> Catalytic cracking produces alkanes, cycloalkanes and aromatics.

> The mechanism involves the formation of carbocations.


16 Joel Biffin

Combustion of Alkanes

- Complete combustion occurs in the presence of oxygen, the combustion of alkanes is highly exothermic

- Incomplete combustion occurs when there is a lack of oxygen present.

- Sulfur impurities can be found in petroleum fraction which, when burnt, produce SO2 & SO3.

- The internal combustion engine produces pollutants:

> CO2, H2O, CO, NOx & Unburnt Hydrocarbons.

- Catalytic converters remove these pollutants by converting CO, NOx & Hydrocarbons to CO2, N2 & H2O.

- Converters have a honeycomb coated layer of Pt/Pd/RH - giving a large surface area.

Definitions

Fractional Distillation A method of separating molecules by making use of their different boiling points

Cracking The breaking of C-C bonds in long chain alkanes to give shorter chain, more useful, more valuable hydrocarbons

Thermal Cracking Energy required for bond breaking in hydrocarbon is provided by heat alone. Up to 900oC at 70atm Forms mostly alkanes

Catalytic Cracking

Energy required for bond breaking in hydrocarbon is provided by heat in the presence of a catalyst (zeolite or aluminosilicate) 450oC Forms branched alkanes, cycloalkanes and aromatic hydrocarbons

Fuel Substance which on combustion produces energy

Fossil Fuel Fuel which has been formed over millions of years and which is a finite resource

Greenhouse Gas Any gas which contributes to global warming by absorption of infra-red radiation by their bonds (polar)

CH4(g) + 2O2(g) -> CO2(g) + 2H2O(l)

CH4(g) + 3/2 O2(g) -> CO(g) + 2H2O(l)

CH4(g) + O2(g) -> C(s) + 2H2O(l)

AS Level Chemistry | Unit 1 Revision Guide

Documents

Transcript of AS Level Chemistry | Unit 1 Revision Guide