Wavelength (lambda)
Distance between two peaks or two roughs (measured in m or nm)
Frequency (v, nu)
Number of waves that pass through a given point . (units S-1, HZ)
Constructive vs deconstructive
Waves interfere, do not need to be the exact same.
E = h*v
h = Plank’s constant (6.626*10^-24 J/S
E = E photon (= E binding + KE)
V = frequency
= hc/wavelength
C = constant (3x10^8 m/s)
Diffraction
Electrons diffract.
Emission of light from atoms
Passage of electricity through gas of atoms causes atoms to emit light.
- Only see 4 frequencies on the detecting screen
- Discrete emission lines can only be explained by quantisation of the energy levels of atoms and molecules
Schrödinger model of the atom
H * Ψ = E * Ψ
H = hamiltonian
E = energy
Ψ = wavelength or orbital = a mathematical function describing the shape of a wave
Ψ^2 = Probability of finding and electron at any point around the nucleus,eaus
Quantum Numbers of Atomic Orbital
- Principle quantum number (n) - size of orbital
- Angular momentum quantum number (l) = (n-1)- shape of orbital
- Magnetic quantum number (ml or m) - orientation of orbital (-l <-> +l)
- Electron spin quantum number (ms or s) = +/- 1/2
Radial probability
Distribution, sum of all Ψ^2
Schrödinger model: Energies of orbitals
HΨ = EΨ
E_n = -2.18x10^-18 J *(1/n^2) -DOES NOT depend on l or m
- Subshells have the same energy (they are degenerate)
E = how stable the orbital is having the probability Ψ
Pauli exclusion principle
No two electrons in an atom can have the same value of all 4 quantum numbers
Factors determining an atom’s energy
- Electron-nucleus attractions
- Electron-electron repulsions - electrons further from the nucleus are shielded from full + nuclear charge by electrons closer to nucleus
Zeff = “effective” nuclear charge
= Z actual - shielding electrons
Ens < Enp < End < Enf
<——– higher zeff
Hund’s Rule
If two or more degenerate orbitals are available, one electron goes into each until they are half full. All half fulled orbitals have the same spin quantum number.
Non-directional solid
Ionic solids form because oppositely charged ions are attracted to each other in all directions.
Electronegativity
The ability of an atom in a molecule to attract electrons toward itself.
Types of Bonding
- Non-polar covalent (En difference = 0) i.e. H2 - Electronically symmetrical
- Polar covalent (En difference <=2.0) i.e HF - partial charges
- Ionic (En difference >=2.0) i.e. LiF - full charges, metal + non-metal
VSEPR Model
- Linear (180º - 2)
- Trigonal Planar (120º -3)
- Tetrahedral (109.5º - 4)
- Trigonal Bipyramidal (90º and 120º - 5 “seesaw”)
- Octahedral (90º - 6)
Dipole Moment
A measure of the separation of charge in a molecule arising from the unequal sharing of electrons in polar bonds.
No dipole moment = non-polar
Permanent molecule = polar
Ion-dipole Interaction
(only for mixtures)
Interaction between fully charges ion and partial charges of a polar molecule
- The energy of attraction increases with the charge of the ion and decreases with the square of the distance between ion and dipole
Dipole-dipole interaction
Polar molecules attract one another when they orient with unlike charges close together, but they repel one another when they orient with like charges together.
Hydrogen-Bonding
Dipole-dipole interactions between H and very electronegative elements N and O. I.e. DNA nucleotide base pairs.
London dispersion forces
Occurs between all molecules, strength depends on size, polarisability.
Pressure
Force exerted per area
P = F/A
Units Pa = N/m^2; kPa = 10^3 Pa
1 atm = 760 mmHg
Kinetic Molecular Theory
- Gases made of tiny particles moving completely randomly
- Total volume of particles very small compared to size of container
- Particles do not interact with each other
- Particle collisions are elastic
- Kinetic Energy increases with temperature
