Chemistry Unit 1

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Chemistry Unit 1
1 Mass spectrometry
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 as follows:
1.1.1 1. The sample is first vaporised so that it becomes easier to do the next step
1.1.1.1 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.
1.1.1.1.1 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.
1.1.1.1.1.1 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
1.1.1.1.1.1.1 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.
1.1.1.1.1.1.1.1 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 whole investigtion.
1.2 Relative atomic mass is calculated by the sum [Relative mass x abundance] divided by the total abundance.
2 Electron Configuration
2.1 We now no longer have shells, we have subshells and orbitals.
2.1.1 The first shell can only hold the S orbital
2.1.1.1 The second shell can hold the S orbital and the P orbital
2.1.1.1.1 The third shell can hold the S orbital, the P orbital and the D orbital
2.1.1.1.1.1 The fourth shell can hold the S, P, D and F orbital.
2.1.1.1.1.1.1 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.
2.1.1.1.1.2 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.
2.1.1.1.2 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.
2.1.1.2 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.
2.2.1.1 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...
2.2.1.1.1 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.
2.2.1.1.1.1 The only exceptions to this are Copper 29 and Chromium 24.
2.2.1.1.1.1.1 Cu29: 1s2, 2s2 2p6, 3s2 3p6, 4s1, 3d10
2.2.1.1.1.1.1.1 As you can see, it's more stable for the elctron to be in the 3d orbital, than in the 4s orbital.
2.2.1.1.1.1.2 Cr24: 1s2, 2s2 2p6, 3s2 3p6, 4s1, 3d5
2.2.1.1.2 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
3.3.1.1 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 remove electrons
3.3.2.1 Genreral trend down a group: decreases because of this determinant i.e. shieldining increases.
3.3.3 Electron Confguration
3.3.3.1 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.
3.3.3.2 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 nuclear attraction.
4 Trend in Melting points graph to know:
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 than Magnesium.
4.1.1.1 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.
4.1.1.1.1 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.
4.1.1.1.1.1 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 Bonding
5.1 Ionic
5.1.1 This is electrostatic forces of attraction between oppostiely charged ions.
5.1.1.1 This is where electrons are are given or taken. It typically occurs between metals and non-metals but it doesnt necessarily have to.
5.1.1.1.1 It forms a giant ionic lattice crystal structure. It has the following properties:
5.1.1.1.1.1 Arranged in a regular lattice
5.1.1.1.1.1.1 Solid at room temperature
5.1.1.1.1.1.1.1 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.
5.1.1.1.1.1.1.1.1 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.
5.1.1.1.1.1.1.1.1.1 They are brittle
5.1.2 Ionic compounds form when there are large differences in electronegativity.
5.2 Covalent
5.2.1 Normal Covalent
5.2.1.1 This is the electrostatic forces of attraction between a shaired pair of electrons between two nuclei. Making them very stable.
5.2.1.1.1 Usually forms between a pair of non-metals.
5.2.1.1.1.1 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.
5.2.1.1.1.1.1 Properties:
5.2.1.1.1.1.1.1 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)
5.2.1.1.1.1.1.1.1 Poor conductors, as they have no free electrons or ions.
5.2.1.1.1.1.1.1.1.1 Usually gases or liquid at room temperature.
5.2.1.2 With covalent bonding, there is little, if any difference in electronegativity.
5.2.2 Dative covalent
5.2.2.1 This is where one atom donates both pairs of electrons to an atom that is electron deficient.
5.2.2.1.1 The atom that is the donator, has a lone pair of electrons. The atom that is the acceptor does not have a filled outer shell.
5.2.2.1.1.1 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 properties.
5.3 Metallic
5.3.1 Electrostatic forces of attraction between cations and a sea of delocalised electrons.
5.3.1.1 Properties:
5.3.1.1.1 They are good conductors of electricity
5.3.1.1.1.1 They are good conductors of heat, because they have free electrons that can carry charge and heat,
5.3.1.1.1.1.1 They're malleable as the layers can slide over eachother
5.3.1.1.1.1.1.1 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.
5.3.1.1.1.1.1.1.1 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 Electronegativity
5.4.1 This is the ability of an atom pull a pair of electrons towards it in a covalent bond.
5.4.1.1 Determinants:
5.4.1.1.1 Nuclear charge: more protons, higher electronegativity as the atoms have a stonger ability to pull electrons towards them
5.4.1.1.1.1 Distance and shielding: As you increase distance and shielding, nuclear attraction gets weak and the ability to pull and electron towards itself weakens.
5.4.1.1.1.1.1 So as you go across a period, electronegativity increases.
5.4.1.1.1.1.1.1 As you go down a group, electronegativity decreases.
5.4.1.1.1.1.1.1.1 Therefore, the most electronegative atoms are found at the top right corner of the periodic table.
5.4.1.1.1.1.1.1.1.1 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.
5.4.2.1 Some atoms are electrongetive enough to pull the shared pair of electrons towards it in a covalent bond.
5.4.2.1.1 This make one side of the molecule more negative than the other side
5.4.2.1.1.1 This is called a polar bond, It is formed when there is a significant different in electronegativity.
5.4.2.1.1.1.1 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'
5.4.2.1.1.1.1.1 The dipole moment represent a polar molecule. The arrow is in the direction of the more negative species.
5.4.2.1.1.1.1.1.1 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.
6.1.1.1 We need to remember that: LPLP repulsion > LPBP repulsion > BPBP repulsion
6.1.1.1.1 If you have 2 BP electrons: Shape: Linear, Bond Angle: 180
6.1.1.1.1.1 3 BP electrons: Shape: Trignal Planar, Bond Angle: 120
6.1.1.1.1.1.1 2 BP electrons, 1 LP electron: Shape:Bent, Bond Angle: 117.5
6.1.1.1.1.1.1.1 4 BP electrons: Shape: Tetrahedral, Bond Angle: 109.5
6.1.1.1.1.1.1.1.1 3 BP electrons, 1 LP electrons Shape: Trigonal Pyramidal, Bond Angle: 107
6.1.1.1.1.1.1.1.1.1 2 BP, 2LP Shape: V-shaped, Bond Angle 104.5
6.1.1.1.1.1.1.1.1.1.1 5 BP Shape: Trigonal Bipyramidal, Bond Angle: 90 and 120
6.1.1.1.1.1.1.1.1.1.1.1 6 BP Shape: Octahedral, Bond Angle: 90
6.1.1.1.1.1.1.1.1.1.1.1.1 4BP, 2LP Shape: Square Planar Bond Angle: 90
7 Structures:
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
7.1.1.1 Strong forces of attraction between oppositely charged ions result in a regular lattice.
7.1.1.1.1 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
7.1.2.1 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.
7.1.2.1.1 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
7.1.3.1 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..
7.1.3.1.1 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
7.1.4.1 This is a form of covalent bonding, however, with macromolecules, the covalent bonds extend throughout the entire structure, making it a huge molecule.
7.1.4.1.1 The many strong covalent bonds mean that they have extremely high melting points because there's so many covalent bonds to break.
7.1.4.1.1.1 You need to know the examples:
7.1.4.1.1.1.1 Diamond
7.1.4.1.1.1.1.1 Extremely hard, the hardest substance known to man.
7.1.4.1.1.1.1.2 Made of Carbon
7.1.4.1.1.1.1.3 Has very very high melting points
7.1.4.1.1.1.1.4 Doesn't conduct electricity
7.1.4.1.1.1.2 Graphite
7.1.4.1.1.1.2.1 Made of Carbon
7.1.4.1.1.1.2.2 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 carbon atom.
7.1.4.1.1.1.2.2.1 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.
7.1.4.1.1.1.2.3 Soft as layers can slide over eachother.
7.1.4.1.1.1.2.4 Can carry charge and conduct electricity due to the fourth delocalised electron
7.1.4.1.1.1.2.5 High melting points due to covalent bonding
7.1.4.1.1.1.3 Silicon
7.1.4.1.1.1.4 Silicon Dioxide
8 Intermolecular Forces
8.1 These are forces that act between covalently bonded molecules. They aren't found between ionic or metallic structures.
8.2 There are 3 types that we need to know:
8.2.1 Hydrogen Bonding
8.2.1.1 This acts between hydrogen and either Nitrogen, Oxygen or Fluorine.
8.2.1.1.1 This is because N, O, F are the most electronegative elements.
8.2.1.1.1.1 It occurs between the lone pair on the N, O or F and the Hydrogen completely
8.2.1.1.1.1.1 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.
8.2.1.1.1.1.1.1 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
8.2.1.2 Structure of Ice:
8.2.1.2.1 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
8.2.2.1 These are intermolecular forces of attraction between dipoles on a polar molecule.
8.2.2.1.1 Two oppositely charged dipoles will form an attraction.
8.2.2.1.1.1 And this will form an intermolecular force.
8.2.2.1.1.1.1 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
8.2.3.1 This occus between any molecule or atom. It works between everything
8.2.3.1.1 It works on the idea that atoms are made of positive and negative charges, even if theyre neutral overall.
8.2.3.1.1.1 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 charge.
8.2.3.1.1.1.1 This will then onduce dipoles on another atom. An attraction will form i.e. a van der Waal IMF will form.
8.2.3.1.1.1.1.1 Van der Waal's are always there, between atoms, molecules etc. They also act in addition to any other IMF that is working.
8.2.3.1.1.1.1.1.1 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.
8.2.3.1.1.1.1.1.1.1 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.
9.1.1.1 They have similar chemical properties.
9.1.1.1.1 Their physical properties gradually increase
9.1.1.1.1.1 They have the same functional group
9.2 Alkanes
9.2.1 Properties
9.2.1.1 They are not solube in water because they are non polar. Only polar substances will dissolve in water.
9.2.1.1.1 They have very low boiling points as they only have van der Waal's IMF's between them.
9.2.1.1.1.1 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.
9.2.1.1.1.1.1 They're not very reactive.
9.2.1.1.1.1.2 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
9.2.2.1 It's from crude oil, which is found under ground.
9.2.2.1.1 You get hydrocarbons with different chain length in crude oil, and these are seprarated by fractional distillation.
9.2.2.1.1.1 This is the separation of the the hydrocarbons according to their boiling points. Each chain length has similar boiling points.
9.2.2.1.1.1.1 In the very top fraction, you get the petroleum gases (>25) C1-4
9.2.2.1.1.1.1.1 Then you get the petrol and gasoline fraction (40-200) C4-12
9.2.2.1.1.1.1.1.1 Then you get the Naptha fraction C8-12
9.2.2.1.1.1.1.1.1.1 Then you get the kerosine and paraffin (200-250) C12-16
9.2.2.1.1.1.1.1.1.1.1 Diesel (250-300) C16-18
9.2.2.2 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.
9.2.2.2.1 Thermal Cracking
9.2.2.2.1.1 Requires a high temperature of around 400-900. Also requires a high pressure of around 7 atmospheres.
9.2.2.2.1.1.1 Strong materials are required to withstand the high pressures. The time spent also needs to be monitered because you dont want really short alkanes.
9.2.2.2.1.1.1.1 Products: Alkenes (to make plastics, polymers), short chain alkenes (fuel).
9.2.2.2.2 Catalytic Cracking
9.2.2.2.2.1 Requires temperatures of around 450. Also requires slight pressures, just over 150 kPa
9.2.2.2.2.1.1 You need a catalyst, called zeolite, which is highly pourous.
9.2.2.2.2.1.1.1 Products: Cycloalkanes (medicines, dyes, drugs)Branched chain alkanes (fuels)
9.3 Isomerism.
9.3.1 Structural Isomerism- having the same molecular formula but different structural formula
9.3.1.1 Functional group isomerism
9.3.1.1.1 This is where molecules have the same molecular formula but a different functional group to eachother
9.3.1.2 Chaim isomerism
9.3.1.2.1 This is where molecules have the same molecular formula but their longest chain of carbon atoms is different.
9.3.1.3 Positional Isomerism
9.3.1.3.1 This is where molecules have the same molecular formula but the functional or side group are in different position
9.4 Combustion
9.4.1 Complete
9.4.1.1 Produces CO2 and H2O
9.4.1.1.1 Carbon Dioxide is a greenhouse gas which contrbutes to global warming
9.4.2 Incomplete
9.4.2.1 Will produce C(s), CO(g)
9.4.2.1.1 Carbon paritculates cause global dimming (photochemical smog) and it's a carcinogen and makes breathing difficult.
9.4.2.1.2 Carbon monoxide is a poisonous gas, it's a respiratory inhibitor.
9.4.2.2 Releases less energy than complete combustion
9.4.3 Other gases that are produced:
9.4.3.1 Sulphur dioxide.
9.4.3.1.1 This contributes to acid rain.
9.4.3.1.2 You can get rid of the sulpur by a process called flu-gas desulphurisation.
9.4.3.1.2.1 The chimneys are lined with calcium oxide. The SO2 reacts with the CaO, to form CaSO2, which is called gypsum and used to make plaster.
9.4.3.2 NOx's
9.4.3.2.1 This cause acid rain, contribute to global dimming and make breathing more difficult.
9.4.3.2.2 These can be removed through catalytic converters. We use the catalyst:
9.4.3.2.2.1 Iridium
9.4.3.2.2.2 Platinum
9.4.3.2.2.3 Rhodium
9.4.3.2.2.4 They convert NOx's to N2 and O2. They convert CO and C to CO2. They convert unburned hydrocarbons to CO2 and H2O