Advanced Higher Chemistry - Unit 1

Matthew Orr
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Matthew Orr
Created by Matthew Orr about 5 years ago
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Advanced Higher Chemistry - Unit 1
1 Electromagnetic Spectrum
1.1 Wavelength is the distance between adjacent crests of a wave
1.2 Frequency is the number of wavelengths that pass a certain point in one second
1.3 Increasing wavelength: Gamma rays, X-rays, Ultra-violet, Visible, Infra-red, Microwaves, Radio waves
1.4 E=Lhc/λ
1.4.1 kJ per mol
1.4.2 L=Avogadro's Constant 6.02*10^23
1.4.3 h=Planck's Constant 6.63*10^-34
1.4.4 c=velocity (3*10^8 for light)
1.4.5 λ=wavelength
2 Oxidation States
2.1 Uncombined elements are 0
2.2 Monatomic ions are the charge on the ion e.g. Cl- is -1
2.3 In most of its compounds, oxygen is -2
2.4 In most of its compounds, hydrogen is +1
2.4.1 In metallic hydrides, it is -1
2.5 Fluorine is always -1
2.6 Oxidation is an increase in oxidation number. Reduction is a decrease in oxidation number.
3 Shapes of Molecules
3.1 Resonance Structures
3.1.1
3.1.2 The bonds in ozone are all the same length
3.2 Shapes
3.2.1 2 electron pairs give a linear arrangement
3.2.1.1 BeCl2
3.2.2 3 electron pairs give a trigonal arrangement
3.2.2.1 BCl3
3.2.3 4 electron pairs give a tetrahedral arrangement
3.2.3.1 CH4
3.2.4 5 electron pairs give a trigonal bipyramidal arrangement
3.2.4.1 PCl5
3.2.5 6 electron pairs give an octahedral arrangement
3.2.5.1 SF6
4 Ionic Lattices, Superconductors and Semiconductors
4.1 Ionic Lattices
4.1.1 Face Centred Cubic
4.1.1.1 NaCl
4.1.1.2 Cation:Anion ratio less than 0.8
4.1.2 Body Centred Cubic
4.1.2.1 CsCl
4.1.2.2 Cation:Anion ratio more than 0.8
4.2 Superconductors
4.2.1 Superconductors have almost no electrical resistance at low temperatures
4.2.2 "High temperature" superconductors can remain superconductive above -196 C
4.2.2.1 Used in transmission of energy, electronics and medicine (MRI)
4.3 Semiconductors
4.3.1 These are part of the metalloids
4.3.2 Electrical conductivity increases with temperature
4.3.2.1 Also increases with exposure to light
4.3.3 These can be "doped" with either a group 3 or group 5 element
4.3.3.1 Group 5 gives an "n" type semiconductor with an extra electron which increases the conductivity
4.3.3.2 Group 3 gives a "p" type semiconductor with a small positive hole which increases the conductivity
5 Chemical Bonding
5.1 Electronegativity
5.1.1 A high difference in electronegativity values between the elements gives an ionic bond while a low one gives a non-polar covalent bond
5.2 Dative Bonds
5.2.1 These occur when one atom provides both electrons for a covalent bond
5.3 Lewis Electron Dot Diagrams
5.3.1 These are used to represent bonding and non-bonding pairs in molecules and polyatomic atoms
5.3.2
6 Orbital Shapes
6.1 There are four main quantum numbers which determine the properties of the atom
6.1.1 Principal=n. This describes the shell occupied by an electron
6.1.2 Angular momentum=l. This describes the shape of the orbital
6.1.3 Magnetic=m. This describes the orientation of the orbital
6.1.4 Spin=s. This describes the electron spin of the electron in the orbital
6.2 Heisenberg's Uncertainty Principle
6.2.1 It is impossible to define with absolute precision, simultaneously, both the position and momentum of an electron
6.3 Pauli Exclusion Principle
6.3.1 No two electrons in an atom can have the same four quantum numbers
6.4 Orbitals
6.4.1 S Orbitals
6.4.1.1 Spherical in shape
6.4.1.2 Only have one orientation
6.4.2 P Orbitals
6.4.2.1 Dumbbell shaped
6.4.2.2 Three orientations: px, py and pz
6.4.2.2.1 px lies on x-axis etc
6.4.3 D Orbitals
6.4.3.1 Four are double dumbbells, one is dumbbell with a ring around the centre
6.4.3.2 Five orientations: dz^2, dx^2-y^2, dxy, dyz, dxz
6.4.3.2.1 dz^2 has a ring, dx^2-y^2 lies on x and y axes, rest lie between corresponding axes
7 Transition Metal Complexes
7.1 Ligands and Complexes
7.1.1 Ligands are negative ions or uncharged molecules with one or more lone pairs of electrons
7.1.1.1 Monodentate ligands have one lone pair
7.1.1.2 Bidentate ligands have two lone pairs
7.1.2 Complexes are metal ions surrounded by ligands
7.2 Naming Complexes
7.2.1 Ligands first in alphabetical order then the metal and its oxidation state
7.2.2 Ligands that end with -ide change the ending to -o
7.2.3 Ammonia becomes ammine and water becomes aqua
7.2.4 If the complex is a negative ion, the metal ends in -ate
7.2.5 Iron becomes ferrate
7.3 Colour in Transition Metal Complexes
7.3.1 The colour of a complex is caused by the splitting of the 3d orbital of the central metal ion
7.3.2 CN->NH3>H2O>OH->F->Cl->Br->I-
7.3.3 The ligands cause the split
7.3.4 The colour seen will be complimentary to the colour that corresponds to the colour with the same energy as that of the d-d transition
8 Oxides, Chlorides and Hydrides
8.1 Oxides
8.1.1 Go from Na2O to Cl2O
8.1.2 Go from ionic lattice to covalent molecular
8.1.3 Go form basic to amphoteric and then to acidic
8.2 Chlorides
8.2.1 Go from NaCl to Cl2
8.2.2 Go from ionic lattice to covalent molecular
8.2.3 The first two are soluble in water, AlCl3 to SCl2 all give off fumes of HCl and Cl2 dissolves to form an acidic solution
8.3 Hydrides
8.3.1 Go from NaH to HCl
8.3.2 Go from ionic lattice to covalent molecular
8.3.3 Go from strongly alkaline to neutral then to strongly acidic
9 Emission Spectra
9.1 Each element produces a different emission spectrum
9.2 The lines seen are due to photons being emitted when electrons drop from one energy level to a lower one. Each line corresponds to a certain energy which in turn corresponds to a colour
9.3 When an electron absorbs energy of a certain amount, it moves from its ground state to a higher energy level.
9.4 There are three main series of lines: the Paschen Series, the Lyman series and the Balmer series
9.4.1 The Paschen series is the infra-red area of the spectrum and involves electrons falling to the n=3 energy level
9.4.2 The Lyman series is the ultra-violet area of the spectrum and corresponds to falls into the n=1 energy level
9.4.3 The Balmer series is the visible area of the spectrum and corresponds to falls into the n=2 energy level
10 Electronic Configuration
10.1 Aufbau Principle
10.1.1 Sub-shells with lower energy are filled first
10.1.2 4s filled before 3p
10.2 Hund's Rule
10.2.1 Electrons fill degenerate orbitals singly before pairing up
10.3 Spectroscopic Notation
10.3.1 e.g. Lithium: Normal= 2, 1 Spectroscopic= 1s^2, 2s^1
10.4 Orbital Box
10.4.1