Unit 1
AQA
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Atomic structure
Amount of substance
Bonding
Periodicity
Introduction to organic chemistry
Alkanes
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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
• 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 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 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 determines the mass of separate atoms and there are 4
main stages
1) 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/watc
h?v=J-wao0O0_qM
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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)
• 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
Ba
Na Mg Al Si P
S Cl Ar
• 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
NaMg Al Si P S Cl Ar
Na
Mg
Al
Si
P
S
Cl
Ar
• 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.
• Relative atomic and molecular masses, the Avogadro 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 percentage yields
• Ar = Atomic mass
• Mr = Molecular mass
n = Moles
average mass of one atom of an element
• The atomic mass is
1/12th mass of an atom of C12
• 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
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PV=nRT
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P=Pressure (Pa)
V=Volume (mᶟ)
n=Moles (m)
R=Constant (8.31)
T=Temperature (K)
Empirical formula= the smallest whole number ratio of each atom of
each element in a molecule
Molecular formula= the actual number of atoms of each element in a
molecule
1. Find the masses of each element
2. 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 formula
4. 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
• Concentration= mol dm⁻ᶟ
• Volume= dmᶟ
• Moles= n
• 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
• % Atom economy =
Mass of desired product
Total mass of reactants
• Yield of a chemical reaction =
x100
Number of moles in a specified product
Theoretical maximum number of moles ofx100
the product
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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
• 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
• 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- the power of an atom to withdraw electron
density in a covalent bond
The arrow represents the increase in
electronegativity
(Cs is the least, and F is the most
electronegative)
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.
= Delta (slightly)
• 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
Van 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
• Solid -> Liquid (Melting/ Enthalpy of Fusion)
• Liquid -> gas (Boiling/ Enthalpy of Vaporisation)
Bonding
• 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
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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
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The periodic table
Trends in the properties of elements of period 3
More trends in the properties of the elements in period 3
A closer look at ionisation energies
• 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.
Melting/boiling point – the energy needed to overcome forces between molecules
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
• Atomic radii decreases because the increased nuclear charge
attracts the electrons closer to the nucleus
• The first ionisation energy generally increases
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)
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
• Carbon compounds
• Nomenclature- naming organic compounds
• Isomerism
• 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
Types 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
1) 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
Structural 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)
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Alkanes
Fractional distillation of crude oil
Industrial cracking
Combustion of alkanes
General formula- Cn H2n+2
Number of carbon atoms in the
longest chain
Name
1
Methane
2
Ethane
3
Propane (three parts on a propeller
Physical properties
4
• Non polar
5
6
• Insoluble in water
• Relatively unreactive
• Boiling point increases with chain length
Butane (four holes in a button)
Pentane (pentagon)
Hexane (hexagon)
• 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.
• 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
Products includeCarbon particles, Nitrogen Oxides, Sulphur Dioxide, Carbon Dioxide,
un-burnt Hydrocarbons and Water vapour
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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