1.1 This is an instrument, or a piece of equipment
that is used to analyse the elements in a
compound. The process by which it occurs is
1.1.1 1. The sample is first vaporised so that it becomes
easier to do the next step
184.108.40.206 2. Ionisation- This is where a beam of electrons are fired at the sample from an
electron gun, and these electrons knock off some of the electrons from the atoms that
have just been vaporised. The electrons usually knock off one, sometimes two,
electrons to produce positive 1+ or 2+ ions.
220.127.116.11.1 3. These positive ions will be attracted towards an electric field (because positive and
negative attract) and the accelerated through two negatively charged plates. The
positive ions will emerge as a fine beam.
18.104.22.168.1.1 4. The fine beam will enter a magnetic field and be deflected depending on their m/z
ratio. This means the lightest ions will be deflected more that those with a higher m/z
ratio. You can change the strength of the magnetic field to deflect theions you want
22.214.171.124.1.1.1 The ions are then directed on to a negative plate. Here, electrons are transferred and a
small current is created which is then amplified on to a computer. The magnitude of the
current depends of the abundance of the isotopes. The group of ions that create the
largest current will have the highest abundance, i.e. there will be more of it.
126.96.36.199.188.8.131.52 Condition required: a Vacuum, so that the air particles don't get in the way
when the ionisation process occurs. This could potentially mess up the
1.2 Relative atomic mass is calculated by the sum
[Relative mass x abundance] divided by the total
2 Electron Configuration
2.1 We now no longer have shells, we have subshells and
2.1.1 The first shell can only hold the S orbital
184.108.40.206 The second shell can hold the S orbital and the P
220.127.116.11.1 The third shell can hold the S orbital, the P orbital and the D
18.104.22.168.1.1 The fourth shell can hold the S, P, D and F
22.214.171.124.1.1.1 The F orbital goes round in groups of 7, so in total, it can hold up to 14
electrons. The fourth shell can hold 32 electrons in total.
126.96.36.199.1.2 The D orbital goes round in groups of 5, so in total, it can hold up to 10
electrons in total. The third shell can hold 18 electrons in total.
188.8.131.52.2 The P orbital goes round in groups of 3, so in total, all the P orbitals can
hold a maximum of 6 electrons (2 x 3) The third shell can hold 8 electrons
in total if we include the fact that it also has an S orbital.
184.108.40.206 S orbital goes round by itself, so the first shell can only
hold up to 2 electrons
2.1.2 Each orbital can have upto 2
electrons per orbital. However,
some orbitals go round in groups.
2.2 Some Rules
2.2.1 1. You always, always start from the lowest
energy level first and then work yourself up
through the shells and orbitals depending on
which orbital the shell allows.
220.127.116.11 2. The only exception to this rule is when you hit the 3rd shell.
Instead of going : 1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2.. You have to
put the 4s orbial before the 3d orbital like so: 1s2, 2s2, 2p6, 3s2,
3p6, 4s2, 3d10...
18.104.22.168.1 The reason for this is because the orbtals go accoridng to energy level order, and
sometimes, the enrgy levels overlap so that actually, the 3d orbital has a higher
energy than the 4s orbital so it comes after.
22.214.171.124.1.1 The only exceptions to this are Copper 29 and Chromium 24.
126.96.36.199.2 Bear in mind that when you remove
electrons, you always remove from the
4s orbital first, then the 3d orbital.
3 Ionisation energies
3.1 This is defined as the minimm mount of energyu
required to remove one mole of electrons from one
mole of gaseous atoms
3.2 Equations: e.g. for sodium (first ionisation energy):
Na (g) -------> Na+ (g) + e-
3.3 Trends and determinants:
3.3.1 Nuclear charge: the more protns there are, the higher the nuclear
charge, the higher or stronger the nuclear attarction, the more
energy required to get rid of the electron
188.8.131.52 Gnereal trend across a period: increases because
of this determinant i.e. nuclear charge increases.
3.3.2 Shielding and Distance: the more shielding, the greater the distance between
them, the weaker the nuclear attraction and so less energy required to
184.108.40.206 Genreral trend down a group: decreases
because of this determinant i.e. shieldining
3.3.3 Electron Confguration
220.127.116.11 High energy subshell: If you're in a higher energy subshell i.e. 3p1
compared to 3s2, then you're more unstable and so not much energy is
required to remove an electron.
18.104.22.168 If you're paired up, then there is repulsion between the two electrons.
Therefore, there is repulsion and so less energy will be required to
get rid of an electron.
3.4 As you remove more and more electrons, more and more energy will be requried.
the reason for this is because when you remove one electron from an atom, you
immediately have more protons than you have electrons, which then means you
have higher nuclear charge each time you remove an electron and so higher
4 Trend in Melting points graph
4.1 Magnesium has a higher melting point than sodium because it has a
higher nuclear charge even though shielding and distance remain the
same, and this means theres a higher nuclear attraction between the
cations and the delocalised electrons so more energy is required than
sodium to break these forces of attraction.
4.1.1 This is the same for why Aluminium has a higher melting point
22.214.171.124 Silicon has a huge melting point because it has macromolecular
structure, so a huge amount of energy will be required to break the
many covalent bonds that silicon has.
126.96.36.199.1 Phosphorus, sulphur and chlorine are simple molecular.
they only have van der Waal's forces between them and
so not much energy is required to break intermolecular
forces that are between them.
188.8.131.52.1.1 Sulphur has a higher melting point than phosphrus
because it goes around as S8 compared to P4, and so
more van der Waal's will form between S8 and so a little
more energy is required to break the IMF's.
5.1.1 This is electrostatic forces of
attraction between oppostiely
184.108.40.206 This is where electrons are are given or taken. It typically occurs
between metals and non-metals but it doesnt necessarily have
220.127.116.11.1 It forms a giant ionic lattice crystal structure. It has the following properties:
18.104.22.168.1.1 Arranged in a regular lattice
22.214.171.124.1.1.1 Solid at room temperature
126.96.36.199.188.8.131.52 High melting point as the fact that there are
electrostatic forces of attraction between the
ions in all directions, loads of energy is
required to overcome these forces of attraction.
184.108.40.206.220.127.116.11.1 They conduct electricity when molten or aqueous as you have
free ions and can carry charge. When solid, the ions are stuck
and vibrate in a fixed position.
18.104.22.168.22.214.171.124.1.1 They are brittle
5.1.2 Ionic compounds form when there are
large differences in electronegativity.
5.2.1 Normal Covalent
126.96.36.199 This is the electrostatic forces of attraction
between a shaired pair of electrons between
two nuclei. Making them very stable.
188.8.131.52.1 Usually forms between a pair of non-metals.
184.108.40.206.1.1 They create molecules even if its a covalent bond between two of
the same atoms. And sometimes, you can also get 2 pairs of
electrons that are shared between two molecules e.g. O2.
220.127.116.11.18.104.22.168 Low melting points. Even though the actual covalent bond is very strong, the
intermolecular forces between the molecules arent strong at all, so noit much
energy is required to break these bonds. (For simple molecular crystals)
22.214.171.124.126.96.36.199.1 Poor conductors, as they have no free electrons or ions.
188.8.131.52.184.108.40.206.1.1 Usually gases or liquid at room temperature.
220.127.116.11 With covalent bonding, there is little, if
any difference in electronegativity.
5.2.2 Dative covalent
18.104.22.168 This is where one atom donates both
pairs of electrons to an atom that is
22.214.171.124.1 The atom that is the donator, has a lone pair of electrons.
The atom that is the acceptor does not have a filled outer
126.96.36.199.1.1 Once the bond has formed, it is impossible to tell
the difference between a dative covalent bond and
a normal covalent bond, they have the exact same
5.3.1 Electrostatic forces of attraction between
cations and a sea of delocalised electrons.
188.8.131.52.1 They are good conductors of electricity
184.108.40.206.1.1 They are good conductors of heat, because they
have free electrons that can carry charge and
220.127.116.11.1.1.1 They're malleable as the layers can slide over eachother
18.104.22.168.22.214.171.124 They're strong. As you go across a period, the nuclear charge
increases, there's a stronger electrostatic force of attraction and so the
metal is stronger.
126.96.36.199.188.8.131.52.1 High melting points. As you go across the period, the melting points tend to
increase because nuclear charge increases and shielding stays the same and
so there's higher nuclear attraction and higher electrostatic forces of attraction,
so more energy will be required to overcome this.
5.4.1 This is the ability of an atom pull a pair
of electrons towards it in a covalent
184.108.40.206.1 Nuclear charge: more protons, higher electronegativity as the atoms
have a stonger ability to pull electrons towards them
220.127.116.11.1.1 Distance and shielding: As you increase distance and shielding, nuclear
attraction gets weak and the ability to pull and electron towards itself
18.104.22.168.1.1.1 So as you go across a period, electronegativity increases.
22.214.171.124.126.96.36.199 As you go down a group, electronegativity decreases.
188.8.131.52.184.108.40.206.1 Therefore, the most electronegative atoms are found at the
top right corner of the periodic table.
220.127.116.11.18.104.22.168.1.1 The most electronegative atoms are Fluorine, Oxygen,
Nitrogen and Chlorine.
5.4.2 The fact that atoms are able to have electronegativity means that
you can have a polar bond.
22.214.171.124 Some atoms are electrongetive enough to pull the shared pair of electrons towards it
in a covalent bond.
126.96.36.199.1 This make one side of the molecule more negative than
the other side
188.8.131.52.1.1 This is called a polar bond, It is formed when there is a significant different in electronegativity.
184.108.40.206.1.1.1 As shown on the diagram below, the chlorine is more
electronegative than the hydrogen, so the electron are
attracted towards the CL, this creates a dipole, which means
a 'small chnage in electron density'
220.127.116.11.18.104.22.168 The dipole moment represent a polar molecule. The
arrow is in the direction of the more negative species.
22.214.171.124.126.96.36.199.1 Species' that are symmetrical may have polar bonds
but dont necessarily need to be a polar molecule.
6 Molecular Shpes
6.1 When a covalent bond is formed, you get a bondiong pair of
electrons, and you may also have lone pairs of electrons. It is
important to understand that electron pairs, regardless of
whether they're bonding pairs or lone pairs, will repel each other.
6.1.1 To minimise repulsion, the electron pairs maximise the space between them i.e.
they will spread out as far as possible.
188.8.131.52 We need to remember that: LPLP repulsion > LPBP repulsion > BPBP repulsion
184.108.40.206.1 If you have 2 BP electrons: Shape: Linear, Bond Angle: 180
220.127.116.11.1.1 3 BP electrons: Shape: Trignal Planar, Bond Angle: 120
18.104.22.168.1.1.1 2 BP electrons, 1 LP electron: Shape:Bent, Bond Angle: 117.5
22.214.171.124.126.96.36.199 4 BP electrons: Shape: Tetrahedral, Bond Angle: 109.5
188.8.131.52.184.108.40.206.1 3 BP electrons, 1 LP electrons Shape: Trigonal Pyramidal, Bond Angle: 107
220.127.116.11.18.104.22.168.1.1 2 BP, 2LP Shape: V-shaped, Bond Angle 104.5
22.214.171.124.126.96.36.199.1.1.1 5 BP Shape: Trigonal Bipyramidal, Bond Angle: 90 and 120
188.8.131.52.184.108.40.206.220.127.116.11 6 BP Shape: Octahedral, Bond Angle: 90
18.104.22.168.22.214.171.124.126.96.36.199.1 4BP, 2LP Shape: Square Planar Bond Angle: 90
7.1 Crystals are solids that are held together by forces of attraction. These
forces of attraction could be Chemical or Physical bonds.
7.1.1 Giant Ionic Crystals
188.8.131.52 Strong forces of attraction between oppositely
charged ions result in a regular lattice.
184.108.40.206.1 They have high melting points and boiling points. The higher the
nuclear charge, the higher the attraction, the more energy needed to
overcome these bonds.
7.1.2 Giant Metallic Crystal
220.127.116.11 Metals also exist as a lattice where a regular repeating
structure of cations are held together by strong
electrostatic forces of attraction betweem the cations
and delocalised electrons.
18.104.22.168.1 They also have high melting points as lots of energy is required to overcome
these forces of attraction. The higher the nuclear charge, the higher the
attraction, the more energy needed to overcome these bonds.
7.1.3 Simple Molecular Crystals
22.214.171.124 These are covalently bonded molecules. They're not
held together chemically, but physically, meaning
through intermolecular forces. IMF's are much weaker
than chemical bonds so much energy isnt required to
overcome these forces of attraction..
126.96.36.199.1 An example you need to know is Iodine. The I2 molecules are held together by
intermolecular forces. Iodine sublimes easily, meaning it turns from a solid to a
gas easily because much energy is not needed to break all the intermolecular
forces between the molecules.
7.1.4 Macromolecular Crystals
188.8.131.52 This is a form of covalent bonding, however, with
macromolecules, the covalent bonds extend throughout the
entire structure, making it a huge molecule.
184.108.40.206.1 The many strong covalent bonds mean that they have extremely
high melting points because there's so many covalent bonds to
220.127.116.11.1.1 You need to know the examples:
18.104.22.168.22.214.171.124 Extremely hard, the hardest substance known to man.
126.96.36.199.188.8.131.52 Made of Carbon
184.108.40.206.220.127.116.11 Has very very high melting points
18.104.22.168.22.214.171.124 Doesn't conduct electricity
126.96.36.199.188.8.131.52 Made of Carbon
184.108.40.206.220.127.116.11 Structure: each carbon is covalently bonded to 3 other carbons (trigonal planar). Carbon
makes 4 bonds, so that must mean there's 1 delocalised electron per
18.104.22.168.22.214.171.124.1 Graphite forms layers of carbon, between each layer are the delocalised
electrons. Van der Waal's also form betwene the layers to hold it together.
126.96.36.199.188.8.131.52 Soft as layers can slide over eachother.
184.108.40.206.220.127.116.11 Can carry charge and conduct electricity due to the
fourth delocalised electron
18.104.22.168.22.214.171.124 High melting points due to covalent bonding
126.96.36.199.1.1.4 Silicon Dioxide
8 Intermolecular Forces
8.1 These are forces that act between covalently bonded
molecules. They aren't found between ionic or metallic
8.2 There are 3 types that we need to know:
8.2.1 Hydrogen Bonding
188.8.131.52 This acts between hydrogen and
either Nitrogen, Oxygen or
184.108.40.206.1 This is because N, O, F are the most
220.127.116.11.1.1 It occurs between the lone pair on the N, O or F and the
18.104.22.168.1.1.1 It's the strongest type of intermolecular
force, so has the highest melting point
because more enrgy will be required to
over come the hydrogen bonds.
22.214.171.124.126.96.36.199 To get the marks in the exam, you need to put the lone pair on the N,O or F.
You also need the delta + on the hydrgoen and 3 dash lines (- - -) directly
between the hydrogen and the lone pair found on the N,O or F
188.8.131.52 Structure of Ice:
184.108.40.206.1 It's less dense than water. The reason for this is because water is
liquid, so the hydrogen are constantly made and broken, however,
ice is a solid and so the hydrogen bonds hold the water molecules
in a fixed position. The regular arrangement means that the water
molecules are more spread out, and so less dense.
8.2.2 Dipole-Dipole Forces
220.127.116.11 These are intermolecular forces of attraction
between dipoles on a polar molecule.
18.104.22.168.1 Two oppositely charged dipoles will form
22.214.171.124.1.1 And this will form an intermolecular force.
126.96.36.199.1.1.1 In the Exam, you need to show the delta + or -
and a dash line betweent the delta + and delta -.
8.2.3 Van der Waals
188.8.131.52 This occus between any molecule or atom. It works between everything
184.108.40.206.1 It works on the idea that atoms are made of positive
and negative charges, even if theyre neutral overall.
220.127.116.11.1.1 Electrons are constantly moving, so at any one time, one side of an atom
could be more negative than another. This then creates a temporary
18.104.22.168.1.1.1 This will then onduce dipoles on another atom. An attraction will form i.e.
a van der Waal IMF will form.
22.214.171.124.126.96.36.199 Van der Waal's are always there, between atoms, molecules etc. They also
act in addition to any other IMF that is working.
188.8.131.52.184.108.40.206.1 The more electrons there are, the bigger the atom and the more van der Waal's formed.
this then means more enrgy will be required to break the IMF.
220.127.116.11.18.104.22.168.1.1 This is why usually, as you go down a group, the melting and boiling points
increase because more energy will be required to break the IMF. It's the same
pattern when you increase the chain length of the Hydrocarbons.
9 Organic Chemistry- The chemistry of carbon atoms
9.1 A homologous series has the following properties:
9.1.1 The chain increases by CH2 each time.
22.214.171.124 They have similar chemical properties.
126.96.36.199.1 Their physical properties gradually increase
188.8.131.52.1.1 They have the same functional group
184.108.40.206 They are not solube in water because they are
non polar. Only polar substances will dissolve
220.127.116.11.1 They have very low boiling points as they only have van der
Waal's IMF's between them.
18.104.22.168.1.1 However, as you go down the group, the boiling points increase
because as you increase the chain number, you increase the
number of electrons, so more van der Waal's are formed.
22.214.171.124.1.1.1 They're not very reactive.
126.96.36.199.1.1.2 Branched chain alkenes tend to have lower boiling points than straight
chained alkanes. The reason for this is because they have an irregular
pattern and so not as many van der Waal's can form, hence a lower
boiling point. This is why they're used as fuels, they ignite easily.
9.2.2 Derivatives and uses
188.8.131.52 It's from crude oil, which is
found under ground.
184.108.40.206.1 You get hydrocarbons with different chain length in
crude oil, and these are seprarated by fractional
220.127.116.11.1.1 This is the separation of the the hydrocarbons according
to their boiling points. Each chain length has similar boiling
18.104.22.168.1.1.1 In the very top fraction, you get the petroleum gases (>25) C1-4
22.214.171.124.126.96.36.199 Then you get the petrol and gasoline fraction (40-200) C4-12
188.8.131.52.184.108.40.206.1 Then you get the Naptha fraction C8-12
220.127.116.11.18.104.22.168.1.1 Then you get the kerosine and paraffin (200-250) C12-16
22.214.171.124.126.96.36.199.1.1.1 Diesel (250-300) C16-18
188.8.131.52 Some alkanes are higher in demand than others. In order
to supply more, we need to crack some of the longer
chained alkanes into shorter chained alkanes.
184.108.40.206.1 Thermal Cracking
220.127.116.11.1.1 Requires a high temperature of around 400-900. Also requires a
high pressure of around 7 atmospheres.
18.104.22.168.1.1.1 Strong materials are required to withstand the
high pressures. The time spent also needs to
be monitered because you dont want really
22.214.171.124.126.96.36.199 Products: Alkenes (to make plastics, polymers), short chain
188.8.131.52.2 Catalytic Cracking
184.108.40.206.2.1 Requires temperatures of around 450. Also requires slight
pressures, just over 150 kPa
220.127.116.11.2.1.1 You need a catalyst, called zeolite, which is highly pourous.