1. Ch. 13 Electrons in Atoms Ch. 13.1 Models of the Atom
Ch Electron Arrangement in Atoms
Ch Physics and the Quantum
Mechanical Model

2. Ch. 13.1 Models of the Atom Evolution of Atomic Models
John Dalton – atomic theory
JJ Thomson – plum pudding model
Ernst Rutherford - nuclear atom
Niels Bohr – planetary model
Fixed energy levels
Quantums of energy
Schrodinger – quantum mechanical model

3. Ch. 13.1 Models of the Atom The quantum mechanical model
Primarily mathematical model
Restricts energy of electrons
Estimates probability if finding an electron in a certain position
90% probability

4. Ch. 13.1 Models of the Atom Atomic orbitals
Designates energy levels of electrons with principal quantum numbers (n)
n=1, n=2, n=3, n=4 and so on
The principal energy levels have a specific number of sublevels
The sublevels are designated s, p, d, and f
Sublevels have specific numbers of orbitals
The orbitals have specific shapes that correspond to the path of the electrons
Nodes are areas where the probability of finding electrons is low

5. Ch. 13.2 Electron Arrangement in Atoms
Electron Configurations
The ways in which electrons are arranged around the nucleus of an atom
Three rules explain how to find the electron configuration of atoms
Aufbau principle – e- enter the lowest energy level first
Pauli exclusion principle – an atomic orbital may hold at most 2 e- (with opposite spin)
Hund’s rule – when e- occupy orbitals of equal energy, one e- enters each orbital until all the orbitals contain one e- with parallel spins
Exceptional electron configurations – Cr, Cu

6. Ch. 13.3 Physics and the Quantum Model
Light and atomic spectra
Electromagnetic radiation
Amplitude – height of wave from origin
Frequency – measured in hertz (Hz)
Wavelength – distance between crests (nm)
Speed – speed of light (a constant, c)
v = c/l
Frequency = speed/wavelength
Visible spectrum (ROY G BIV)
Atomic emission spectra (fingerprint)

7. Ch. 13.3 Physics and the Quantum Model
The quantum concept and the photoelectric effect
Max Planck
Amount of energy absorbed or emitted is proportional to the frequency of the radiation
E = hv (energy = Planck’s constant x frequency)
h = x Js
Albert Einstein
Light can be described as quanta of energy (photons)
Dual nature of light (waves and particles)
First to explain the photoelectric effect
Energy must meet a threshold value to eject e- from the surface of metal

8. Ch. 13.3 Physics and the Quantum Model
An explanation of atomic spectra
When energized, e- move from the ground state to an excited state (n = 2,3,4, etc)
The same amount of energy that was absorbed is then emitted as light and the e- falls back down to a lower energy level
Results in Lyman, Balmer and Paschen series
An upper limit for the frequency of light emitted exists because a very excited e- will escape the atom
Bohr’s model of the atom was not able to explain bonding and was eventually replaced

9. Ch. 13.3 Physics and the Quantum Model
Quantum mechanics
De Broglie’s equation predicts that all matter exhibits wave-like motions
True for all matter, but motion depends on the size of the object
Wavelengths for objects visible to the naked eye are not measurable
Wavelengths for extremely small objects are measurable
Heisenberg’s uncertainty principle
It is impossible to know exactly both the position and velocity of a particle at the same time
The more precise the measured velocity, the less precise the position, and vice versa