Key Autumn semester Flashcards
Equation for wave length
λv = c
λ Wavelength (m)
v frequency (Hz or s^-1)
c speed of light (2.998 x 10^8 m/s)
Equation for energy
E = hv
E energy (j)
h plancks (6.63 x 106-34 JS)
De Broglie relationship
Energies calculated for a wave and for a particle must be equal for anything behaving as both it shows wave-particle duality
λ = h/p
p momentum of particle (mass x velocity)
Uncertainty principle
You can not know both the position and speed of an electron but can estimate the probability that an electron is in a certain place
Heisebery uncertainty
The more we know the position of the electron the less we know about the momentum
ΔpΔx ≥ h / 4π
Δpx = mass x velocity
Δx = mass
Schrodinger equation
Tells use the likelihood of finding an electron an a particular location and the allowed wave functions
Quantum numbers and allowed values
n (principle quantum number) = which orbital the electron is located (1,2,3,4,5…)
l (orbital angular momentum) = shape of the orbital (0 to (n-1) ) 0 = S, 1 = p, 2 = d and 3 = f
ml (angular quantum number) = orientation of orbital (-l to +l)
Radial wave function
The probability of finding an electron in a spherical shell as r form the nucleus.
Found using n and small chance of finding e- far from the nucleus
Angular wave funtion
Shows the shape and orientation of orbitals
Electronegativity
The ability of an atomic nuclei to attract electron density shown by the Pauline scale
Lewis structures
Shows the connectivity and formal charges on atom in a structure
Resonance
The blending of two or more lewis structures, where the electrons are delocalised resulting in a resonance hybrid which is a lower energy
VSEPR
Valence shell electron pair repulsion which is used to determent he 3d structure of a molecule.
Lone pairs
Located closer to the central atom so has increased repulsion and so counted as a region of electron density and lower bonds angles
Dispersion forces
exist for all molecules and increase with an increase of electron, same as instantaneous induced dipole
Symmetry operation
An action that transforms a molecule so the resultant molecule is complete unchanged
Rotation
this is a rotation around a specific axis given as Cn for a rotation 1/n
C2 = 180 degree rotation (1/2)
C3 = 120 degree rotation (1/3)
C4 = 90 degree rotation (1/4)
C6 = 60 degree rotation (1/6)
Inversion
Centre of inversion (i) If any point of the molecule passes through a centre of inversion an dit remains the same
reflection
When reflection (σ) on a mirror plane either vertical or horizontal leave the molecule exactly the same
rotation-reflection
Sn > Rotating a molecule by Cn around an axis then reflecting it onto a place perpendicular to the same axis
identity
E > Doing nothing to change the molecular at all or 360 degree rotation
Point group
A type of molecule identifier using its symmetry elements it helps determine polarity (cannot be polar if has a centre of inversion, D groups, cubic groups and icosahedral groups) and chiral (can not be chiral is it contains an improper axis of rotation)
Aufbau principle
Lower energy subshells fill first
Pauli principle
No same electron in an atom can have the same set of quantum numbers
allotropy
property of elements being able to exist in two or more different forms. They from in changing conditions, temperature and methods of formation
What is an oxidation free energy diagram/ frost diagram
A picture representation of the relative stabilities of different oxidation groups
redox couple
relative stability of two oxidation states given by free energy changes. If E is negative for reduction the reaction will happen and the product is more stable.
Nernst equation
∆G (reduction) = -nFE (reduction)
rearranged to give
∆G (oxidation)/F = nE
Plotting a oxidation free energy diagram
Plot ∆G(oxidation)/F on the Y axis with the oxidation state on the x axis
∆G/F (oxidation) if worked out by n x E
Interpreting oxidation free energy diagrams
- change down a slope is favourable
- lower/more negative the oxidation sate is the more stable it is
- the steeper the incline the stronger oxidising agent and the steeper the decline the stronger the reducing agent
Disproportionation reactions
same element is both oxidised and reduced in a reaction
comproportionation reaction
reaction of two element with different oxidation states react to form a compound with an intermediate oxidation number (opposite of disproportionation)
Flame test
If an atom or ion is excited an electron can be promoted to a higher energy levels and then fall back down releasing energy as different coloured light dependant on amount of energy.
Light energy to colour
(lowest energy) R - O - Y -G -B -I -V (Highest energy)
Electron sheilding
Blocking of valence shell electrons attraction to the nucleus due to the presence of inner shell electron creating repulsion effects
Slater rules and equation
Zeff = Z - S
(1s)(2s2p)(3s3p)(3d)(4s4p)(4d)(4f)(5s5p)(5d)(5f)
- no contribution to the right of the one being considered
- 0.35 added for each electron in the same group as the one being considered (0.30 if 1s)
- ns ore np orbital then all electron n-1 contribute 0.85 and n-2 contribute (1)
- if nd or nf then all orbital below contribute 1.00
Reason for trend in atomic and ionic radii
- nuclear charge
- electron sheilding
- Zeff
D-block contraction
With transition metals across the group the atomic radius only decreases very slowly this is because same number of s electrons by differing d electrons. The d electron have poor shielding.
The Lanthanide contraction
Poor shielding effects of the 4f electrons increasing attraction to outer orbitals (increased Zeff)
Lattice energy
The measure of the strength of the force between ions in an ionic solid. The greater the lattice enthalpy the stronger the forces.
Kapustinkiis equation
Calculation of lattice enthalpy for nay compound (UL)
UL (kj/mol) = (121400 x (Z+ x Z-) x V / (r+ + r-) x (1-34.5)/(r+ + r-)
Z = ion charge
v = number of ions per formal unit
r - radius of ions (pm)
Effects on lattice energy
lattice energy with increase when
- ion charge increases (Z)
- ion number increases (v)
- ionic radii (r) decreases
Close packing system
the most tightly packed composition of crystal structures, minimising the unfilled space between ionic spheres. It takes into consideration radius ration and cation to anion ratio
Radius ratio
Anion form a structure and cation fill the holes left over. The smaller the radius of ions the closer they can be packed together and this increasing the lattice enthalpy. Smaller cation occupy tetrahedral holes and larger ion octahedral holes. if the cation is to large the anions will adopt a more open structure (cubic array)
Tetrahedral and octahedral sites
tetrahedral > small hole found between 3 anion in one plane and 1 adjacent anions
octahedral > larger hole between 6 anions in an octahedral shape
Oxides, Peroxide and superoxides
oxides = M2O
superoxides = MO2
peroxides = M2O2
Organometallics
a compound containing at least one metal to carbon bond. Smaller cations are most likely to form these as they have the highest charge density.
Crown ether, crytands and host guest complexes
These are all mutidente ligands they are the most stable as there forms allow them to form complexes with metals with low charge densities (group 1 metals) due to their cavity being tailored to fit each cations (maximum attraction)
Covalent to ionic transions
Ionic and covalent bonds are the two extreme of bonding, polar covalent in the intermediate between the two. Polarity (measure of charge separation) is a spectrum where each compound lies
Charge density
The higher the charge density of each ion the stronger the eltrostic attractions and thus the bond
DO ALL CARDS FROM 133 TO 144
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