AS ChemistryUnit 1AQA

Chapters

1. Atomic structure2. Amount of substance3. Bonding 4. Periodicity5. Introduction to organic chemistry6. Alkanes

Chapter 1- Atomic Structure• Fundamental Particles• The arrangement of electrons• Mass number, atomic number and isotopes• The mass spectrometer• More about electron arrangement in atoms• Electron arrangements and ionisation energy

Fundamental Particles• Atoms are made up of three fundamental

particles. Protons, Neutrons and Electrons

Particle Relative charge

Relative mass

Position

Proton + 1 In the nucleus

Neutron None 1 In the nucleus

Electron - 1/1840 Outside the nucleus

The arrangement of electrons• The first shell holds up to 2 electrons• The second shell holds up to 8 electrons• The third shell holds up to 18 electrons

• The electron arrangement for Argon can be written as 2.8.8

Mass number, atomic number and isotopes

• Mass number is always the biggest number because it is the total number of protons and neutrons. In Argon, it is 39.95.

• Atomic number is just the number of protons and is the other number. In Argon, it is 18.

• Isotopes are atoms of the same element with the same number of protons but a different number of neutrons

The mass spectrometer

The mass spectrometer determines the mass of separate atoms and there are 4 main stages1) Ionisation- A beam of electrons is fired at the particles to remove

one or more electrons from the atom’s outer shell creating a positive ion

2) Acceleration- The positive ions are attracted to the negatively charged plates. This accelerates the particles. Lighter ions go faster.

3) Deflection- the beam of particles move into a magnetic field at right angles and are deflected. Heavier ions are deflected less than lighter ions.

4) Detection- the magnetic field is gradually increased so that the separate masses enter the detector at different times. When the charged particles hit the detector, an electric current is made. The greater the current made, the higher the amount of that particular particle is in the sample

http://www.youtube.com/watch?v=J-wao0O0_qM

More about electron arrangement in atoms• There are subshells within an atom.• S subshells can hold 2 electrons• P subshells can hold 6 electrons• D subshells can hold 10 electrons

• In the 1st shell, there is an S subshell (1 subshell in 1st shell)

• In the 2nd shell, there is as S and a P subshell (2 subshells in 2nd shell)

• In the 3rd shell, there is an S, P and D subshell (3 subshells in 3rd shell)

Electron arrangements and ionisation energies• Ionisation energy- the energy required to remove

one electron from each atom in one mole of gaseous atoms.

• Ionisation energies are affected by atomic radius, shielding and nuclear charge

• Ionisation energies decrease down a group (d and d)• Ionisation energies generally increase across a

period (two vowels)

Be Mg Ca Sr BaNa Mg Al Si P S Cl Ar

Electron arrangements and ionisation energies- Anomalies

• In aluminium, the valence electron being removed is from the 3p subshell, which is further away from the nucleus (smaller atomic radius) so it will take less energy to remove than from a 3s subshell

Na Al P Cl

Electron arrangements and ionisation energies- Anomalies

Na Al P Cl

• In Sulphur, the electron being removed is sharing the subshell with another electron, as a result, they will repel each other, making it easier to remove.

Chapter 2- Amount of substance• Relative atomic and molecular masses, the Avoga

dro constant and the mole• The ideal gas equation• Empirical and molecular formulae• Moles in solutions• Balanced equations and reacted calculations• Balanced equations, atom economies and percent

age yields

Relative atomic and molecular masses, the Avogadro constant and the mole

• Ar = Atomic mass• Mr = Molecular mass•

.n = Moles• The atomic mass is• The Avogadro constant (6.022 x10²ᶟ) is the number of

particles in one mole• Avogadro constant = the number of atoms in 12g of

carbon-12

average mass of one atom of an element 1/12th mass of an atom of C12

The ideal gas equation

PV=nRT• P=Pressure (Pa)• V=Volume (mᶟ)• n=Moles (m)• R=Constant (8.31)• T=Temperature (K)

Empirical and molecular formulaeEmpirical formula= the smallest whole number ratio of each atom of each element in a moleculeMolecular formula= the actual number of atoms of each element in a molecule

1. Find the masses of each element2. Find the number of moles of each element (Mass/Mr)3. Find the simplest ratio by dividing by the smallest number

To find the molecular formula4. Divide the total number of atoms in the empirical formula by the Mr of the total molecule (which you will be given)5. Times the empirical formula by the answer

Moles in solutions• Concentration= mol dm⁻ᶟ• Volume= dmᶟ• Moles= n

Balanced equations and related calculations• You must have the same number of atoms on

each side of the equation• Balanced equations tell us the amount of atoms in

a reaction

• In ionic equations, you will have a positive metal and a negative non-metal as products

Balanced equations, atom economies and percentage yields

• % Atom economy = x100

• Yield of a chemical reaction = x100

Mass of desired product Total mass of reactants

Number of moles in a specified product

Theoretical maximum number of moles of the product

Chapter 3- Bonding• The nature of ionic bonding• Covalent bonding• Electronegativity- bond polarity in covalent bonds• Metallic bonding• Forces acting between molecules• States of matter: gases, liquids and solids• Bonding and structure- summary• The shapes of molecules and ions

The nature of ionic bonding• Happens between a metal and a non metal• Electrons are donated from the metal to the non metal so;• The metal becomes positive and the non metal becomes

negative

Properties• Not malleable because a small displacement causes contact

between ions of the same charge causing the structure to shatter.

• Conduct electricity only when molten or in an aqueous solution because only then are the charged particles free to move and carry charge

• Giant structures, so solid at room temperature

Covalent bonding• Between two non metals• Electrons are shared between atoms to create a

stable noble gas arrangement• A single bond is when two electrons are shared• A double bond is when four electrons are shared• Dative covalent ( or coordinate bonding) is when

an atom only receives electrons, and does not share any of its own

• Dative covalent bonds are represented in diagrams by an arrow

Electronegativity- bond polarity in covalent bonds

• Electronegativity- the power of an atom to withdraw electron density in a covalent bond

Electronegativity depends on • The nuclear charge ( greater charge = greater electronegativity)• The distance between the nucleus and the valence electrons (smaller distance = greater

electronegativity)• Shielding of the nuclear charge ( less shielding = greater electronegativity)

• Fluorine is the most electronegative atom• If an atom in a covalent bond attracts electrons more

than the other atom, it will become delta negative, and as a result, the other atom will become delta positive. This makes the molecule polar.

The arrow represents the increase in electronegativity (Cs is the least, and F is the most electronegative)

= Delta (slightly)

Metallic bonding• Between two metals.• Positive metal ions in a sea of delocalised electrons

Properties• Conducts electricity because delocalised electrons can carry a

charge• Conductors of heat because of the vibrations of the closely

packed ions• Strong because of the strong electrostatic attraction between

metal ions and the sea of delocalised electrons• Malleable and ductile because after a small distortion, each

metal ion in in exactly the same environment• High melting points because of the giant structure and the

strong electrostatic attractions

Forces acting between moleculesVan der Waals (weakest)• Instantaneous dipole induced dipole interactions• Acts between all atoms all the time• The dipole is caused by the always changing position of the

electron cloud. More electrons = stronger dipole

Dipole-Dipole• Permanently polar molecules attract because of the

permanent polar charges

Hydrogen Bonds(Strongest)• Special type of Dipole- Dipole interaction ( not an actual bond)• Between Hydrogen and either Oxygen, Nitrogen and Fluorine

States of matter: gases, liquids and solids• Solid -> Liquid (Melting/ Enthalpy of Fusion)• Liquid -> gas (Boiling/ Enthalpy of Vaporisation)

Bonding and structure- summaryBonding• Covalent (Shared electrons. Only non metals)• Ionic (Transferred electrons. Metal and non metal)• Metallic (Metal ions in a sea of electrons. Metal only)

Structure• Simple molecular- Covalent bonds. E.g. CO2

• Macromolecular- Covalent bonds. E.g. Diamond• Giant ionic- ionic bonds• Metallic- Metallic bonds

Intermolecular forces• Weak Van der Waals• Dipole-Dipole• Strong Hydrogen Bonds

Shapes of molecules and ions• One pair of electrons will repel another pair of electrons• These pairs will take positions as far away from each other as possible

Two pairs of electrons• Linear• 180º

Three pairs of electrons• Trigonal planar• 120º

Four pairs of electrons• Tetrahedral• 109.5º

Five pairs of electrons• Trigonal bipyramid• 90º and 120º

Six pairs of electrons• Octahedral• 90º

Chapter 4- Periodicity• The periodic table• Trends in the properties of elements of period 3• More trends in the properties of the elements in p

eriod 3• A closer look at ionisation energies

The periodic table

• All the elements with their highest electron in the s orbital are in the s block

• All the elements with their highest electron in the p orbital are in the p block

• All the elements with their highest electron in the d orbital are in the d block

• Groups go down (Groups Droop down). They have the same number of electrons in their outer shell

• Periods go across.

Trends in the properties of elements of period 3

Na-Mg-Al• Increased charge on ion• Increased delocalised electrons• Increased electrostatic attraction• Therefore- increased melting/boiling point

Si• Have to overcome very strong covalent

bonds throughout the structure• Macro-molecular structure

S₈ > P₄ > Cl₂- > Ar• Bigger molecule = more electrons = more Van der Waals forces• Low Mp/Bp because only weak Van der Waals forces

have to be overcome

Melting/boiling point – the energy needed to overcome forces between molecules

More trends in the properties of elements of period 3• Atomic radii decreases because the increased

nuclear charge attracts the electrons closer to the nucleus

• The first ionisation energy generally increases

A closer look at ionisation energies

Aluminium• has its outer electron

in a 3p subshell, which takes less energy to remove than an electron in the 3s subshell (Mg), so the ionisation energy is lower than Mg

Sulphur• has a spin shares pair

of electrons in its 3p subshell, so the electrons repel each other so they are easier to remove, so the ionisation energy is lower than P. (Sulphur Shares)

A closer look at ionisation energies

Successive ionisation energies

• There are 7 points before the jump, so the element is in group 7

• There are 4 points before the jump, so the element is in group 4

Chapter 5- introduction to organic chemistry• Carbon compounds• Nomenclature- naming organic compounds• Isomerism

Carbon compounds• Carbon has 4 electrons in its outer shell, so it can

form 4 covalent bonds• Carbon-carbon bonds are relatively strong• Carbon can form single, double and triple bonds

Molecular formula- • the formula that shows the actual number of

atoms of each element in the molecule• Can be worked out from the empirical formula and

the Mr of the empirical and molecular formula

Carbon compoundsTypes of formulae;

• Display- shows every bond and atom in the molecule

• Structural- shows the arrangement of atoms in a molecule without showing any bonds – CH3 CH2 CH2 CH2 CH2 CH3

• Empirical- the lowest whole number ratio of each atom of each element in a molecule – C3 H8

• Molecular- the formula that shows the actual number of atoms of each element in the molecule- C6 H16

Nomenclature- naming organic compounds

Nomenclature- naming organic compounds1) Work out how many carbon atoms are in the

longest chain (5)2) Number the carbon atoms and look to see if

there are any side chains and on what carbon atom they are on (2 and 2) (Always go for the lower number when you have a choice between two options)

3) Count the number of carbon atoms on the side chain (one on each)

Groups are written in alphabetical order

2.2 dimethyl-pentane

IsomerismStructural isomerism• Same molecular formula• Different structural formula

Positional isomerism• The functional group is attached to the main chain at

different points

Functional isomerism• When there are different functional groups

Chain isomerism• Hydrocarbons are arranged differently (in chains)

Chapter 6- Alkanes• Alkanes• Fractional distillation of crude oil• Industrial cracking• Combustion of alkanes

Alkanes

General formula- Cn H2n+2

Physical properties• Non polar• Insoluble in water• Relatively unreactive• Boiling point increases with chain length

Number of carbon atoms in the longest chain

Name

1 Methane

2 Ethane

3 Propane (three parts on a propeller

4 Butane (four holes in a button)

5 Pentane (pentagon)

6 Hexane (hexagon)

Fractional distillation of crude oil• Physical process. No

covalent bonds are broken. Only intermolecular bonds are broken.

• Crude oil is separated into fractions of hydrocarbons of a similar boiling point.

Industrial cracking• Long chain alkanes are cracked to produce shorter

chain alkanes and alkenes• This is done because shorter chain hydrocarbons

are in more demand

Thermal cracking• Very high temperature and pressure

Catalytic cracking• High temperature and pressure• Zeolite catalyst

Combustion of alkanesProducts include- Carbon particles, Nitrogen Oxides, Sulphur Dioxide, Carbon Dioxide, un-burnt Hydrocarbons and Water vapour

• Combustion of alkanes produces CO2 and H2O• Incomplete combustion produces CO and H2O• N2 O2 » 2NO• Sulphur dioxide reacts with water in the air to make acid rain• Calcium oxide or limestone is used to remove sulphur dioxide

Catalytic converters• Uses platinum in a honeycomb shape to increase surface

area• 2CO + 2NO » N2 + 2CO2

Global warming• CO2 and water vapour trap infra red radiation in the Earth’s

atmosphere, causing the Earth to heat up