Chem H Unit 6 Quantum Chemistry

Wave

transfer of energy through a disturbance through space and time

Wavelength

Distance between identical points of successive waves

Amplitude(A)

Measure of waves displacement

Intensity(I)

Energy passing through per unit per time(A²)

Frequency(f)

number of waves passing a point per second

Wavelength formula

v=λf

Wave equation for EM radiation:

c=λf

Order of EM spectrum from high frequency to low frequency

Gamma, X, UV, visible, IR, microwave, radio(FM to AM)

Plancks relation

E=hf; graph is a slope

Photoelectric effect

Electrons are ejected from the surface of certain materials when they absorb light of specific frequencies(called threshold frequency)

Problem with Photoelectric effect

Could not be explained by wave theory

Wave-particle duality:

Light acts as both a wave and a particle

de Broglie wave equation

E=hv/λ; λ=h/mv (m is in kg!)

Problems with Rutherford’s model

1. Orbiting electrons are not stable. They are constantly accelarating (rotation), releasing energy as EM radiation in the process. The loss of energy means that the electrons will eventually crash into the nucleus.

2. Could not explain the line spectra: Why was light only emitted/absorbed at specific wavelengths?

3. Could not explain periodic trend of elements

Bohr model

Electrons orbit the nucleus at specific distances(quantized distances) called n

Rydberg constant

2.18×10⁻¹⁸J

Quantum jumps

Atom absorbing/releasing energy causes the electron to jump/fall; more energy→towards violet light

Spectral series

1. Lyman, UV

2. Balmer, UV/vis

3. Paschen, IR

4. Brackett, IR

Problem with Bohr model

Only matched spectra lines of hydrogen and one electron species

Schrodingers Equation

Describes waves. Electrons were waves with specific energies and spacial distributions(quantized). Wavefunction/orbital represented by greek letter psi. Solutions can gbe approximated for atoms with more than one electron

Heisenberg uncertainty principle

It is impossible to know both the position and the momentum of a particle with certainty

Quantum numbers:

n, l, ml, mn

n

shell, orbital size and distance of electron to nucleus

l

subshell, shape of orbital

ml

Orientation of the orbital

ms

spin up/spin down

Stern Gerlach experiment

Predict existence of spin by passing beam of silver atoms through magnetic field. Half deflect up, half deflect down.

Hydrogen orbital energy

All orbitals with the same n have the same energy(degenerate)

Aufbau principle

Lowest energy levels filled before higher energy levels(reflected on periodic table)

Pauli exclusion principle

Electrons in same orbital have opposite spins

Hund’s rule

MMost stable arrangement of electrons are ones with greatest number of parallel spins. Arrows face upwards

Electron configuration exceptions for full d shells:

Copper, gold, silver

Electron configuration exceptions for half filled d shells

Molybdenum, Chromium

Valence electrons composition

number of s and p electrons of outermost occupied shell. Based on periodic table group number

duet rule:

H and He have full outer shell with 2 valence electrons

Coulombs law

1. Opposite charges attract, like charges repel

2. Force increases with increasing charge

3. Force decreases with increasing distance between charges

Zeff

Effective nuclear charge: essentially what charge valence electrons feel since they are shielded from nucleus. Total charge(# protons)-number of core electrons

Atomic radius

Size of an atom determined by size of electron cloud. Half distance between two nuclei in adjacent atoms.

Atomic radius trends

Increases going down and left

Increases going down because of n (shell) increasing

Increases going left because there is a decreasing Zeff, electrons are sucked in less

Ionic radius

Size of ions compared to atoms

Ionic radius Trends:

As atoms become more negatively charged, they become bigger. More electrons means more shielding

Ionization energy (IE)

Minimum energy it takes to remove an electron

Ionization energy trends

Increase ionization going up and right; again bc of Zeff and n levels

Ionization energy trends exceptions:

group 2vs13 and 15vs16. Group 2 has full s shell. Group 15 has half filled shell

Electron affinity (EA)

Energy released when electron is added to atom. How badly an atom wants an electron. More positive →more easily atom excepts. Negative of energy change

EA Trends:

Increases going up and right. Group 17 and 1 have higher EA’s since they are close to having full shells. Group 2 and 18 have lower EAs since their shells are already full.

Metallic character

increases going down and left. Francium is the most metallic since it has the lowest IE.

Nonmetallic character

Increase going up and right. Fluorine is the most nonmetallic since it has high EA.

Diamagnetic

electrons are all paired

Paramagnetic

at least one unpaired electron

Filled d subshells prefered

Au, Ag, Cu

Half filled d subshells prefered

Mo, Cr

Electronegativity

tendency for atom to attract electrons