Electrons in Atoms
Chemistry
Chapter 5
Models of the Atom
 Rutherford’s model of the atom put the
protons & neutrons in the nucleus with
electrons orbit around the nucleus like
planets around the sun.
 This
model (Rutherford’s) did not adequately
explain:
• Why metals give off colors when heated
(photoelectric effect), or
• The chemical properties of elements
The Bohr Model
 Neils Bohr suggested improvements to
Rutherford’s model.
 Bohr
proposed that an electron is found only in
specific circular paths, or orbits, around the
nucleus.
 Each possible orbit had a fixed energy. The fixed
energies an electron can have are called energy
levels.
 Electrons could move from energy level to another if
they gained or lost the right amount of energy
(quantum).
Quantum Leap
 The amount of energy required to move an
electron is not constant.
 Energy
 Bohr’
levels are not evenly spaced;
model worked well for the simplest atoms,
however it still failed in many ways to explain the
energies absorbed & emitted by atoms with multiple
electrons.
The Quantum Mechanical Model
 The quantum mechanical model is a
mathematical equation that describes the
behavior of electrons; based on work done by
Erwin Schrodinger (1926).
 Varies
from the Bohr model in that it does not
restrict an electron to an exact path around the
nucleus.
The quantum mech. model determines the
allowed energies an electron can have and how
likely
it is to find the electron in various
locations around the nucleus.
Bohr Model VS. Quantum
Mechanical Model
 The link below will take you to an
animation comparing Bohr’s atomic
model and the Quantum Mechanical
Mode.
Energy Levels & Atomic Orbitals
(handout)
 Schrodinger’s equation gives the energies an
electron can have; we call these energy levels.
 For each energy level, Schrodinger’s equation
ultimately leads to a mathematical expression,
called atomic orbitals
 An
atomic orbital is often thought of as a region of
space in which there is a high probability of finding
an electron.
Energy Levels & Atomic Orbitals cont.
 Energy levels are labeled by principal
quantum #s (n).
 These
are assigned values
n = 1, 2, 3, 4, etc.
 Each energy level can have mult. sublevels
 n = n sublevels
 For each sublevel there may be several
orbitals with different shapes
n
has n2 orbitals
 Each orbital can hold only 2 electrons
 n can hold 2n2 electrons
More on Atomic Orbitals
 Different atomic orbitals are denoted by letters.
 s orbitals are spherical; only one s orbital
 p orbitals are dumb-bell shaped; three p orbitals
 d orbitals have clover leaf shapes; 5 d orbitals
 f orbitals more complex; 7 f orbitals
Orbitals cont.
 The # and kind of atomic orbitals depend on the
energy sublevel.




The lowest principal E level (n=1) has on one sublevel
called 1s (only 1 orbital)
The 2nd Principal E level (n=2) has 2 sublevels 2s & 2p (4
orbitals)
The 3rd principal E level (n=3) has 3sublevels 3s, 3p, & 3d.
(9 orbitals)
The 4th principal E level (n=4) has 4 sublevels, 4s, 4p, 4d,
& 4f. (16 orbitals)
 Each orbital can hold no more than 2 electrons!!
Electron Configurations
 The way in which electrons are arranged in
various orbitals around the nuclei of atoms are
called electron configurations.
Three rules – the Aufbau principle, the Pauli
exclusion principle, and Hund’s rule – tell you
how to find the electron configurations of atoms.
Aufbau Principle
 According to the Aufbau Principle, electrons
occupy the orbitals of lowest energy first.
Pauli Exclusion Principle
 According to the Pauli exclusion
principle, an atomic orbital may
describe at most two electrons.

For 2 electrons to occupy the same orbital
they must have opposite spins
Hund’s Rule
 When using the Aufbau diagram to decide how
electrons occupy orbitals of equal energy, one
electron enters each orbital (of a sublevel) until all oribtals (in a sub-level)
contain one electron with the same spin.
Hund’s Rule cont.
 For example, electrons would occupy p
orbitals (of equal energy) as follows.
Electron Configurations
 Click on the button below to be directed
to a interactive site on writing electron
configurations.
Writing out electron configurations
Exceptional Electron Configurations
 Some actual electron configurations differ from
those assigned using the aufbau principle
because half-filled sublevels are not as
stable as filled sublevels, but they are more
stable than other configurations.
 Exceptions
include the elements in the two
groups headed by Cr and Cu.
Physics & the Quantum Mechanical
Model
 The quantum mechanical model grew
from the study of light.
 Light
consists of waves.
 There are many different kinds of waves,
however they all share certain basic
characteristics.
 Amplitude
 Wavelength
 Frequency
Click Here
Amplitude
 The maximum (or greatest) movement
from rest is called the amplitude of the
wave.
Amplitude
Normal
resting
position
Wavelength
 The distance between two consecutive (one
after the other) crests or troughs of a wave is
called the wavelength.
wavelength
The symbol for
wavelength is the Greek
letter lambda.
wavelength
Frequency
 The number of complete waves (cycles),
per unit of time (usually seconds) is called
the frequency.

Unit used to measure wave frequency is hertz (Hz).
• May also be expressed as a reciprocal second (s-1).
1st
…50th
1 sec.
What is the frequency
of this wave?
Remember: Frequency = cycles per second = Hertz (Hz)
Frequency represented by v (Greek letter nu).
Light Speed
 The wavelength and frequency of light are
inversely proportional to each other.
 The
product of frequency and wavelength always
equal a constant (c) the speed of light.
C=
v
All electromagnetic waves travel in a
vacuum at the speed of 3.00 x 108 m/s.
All electromagnetic waves travel
in a vacuum at the speed of
3.00 x 108 m/s.
Electromagnetic Spectrum
 The Electromagnetic
spectrum ranges from
very long EM waves
(low frequency radio) to
very short waves (high
frequency gamma
rays).
 The amount of energy
carried by EM waves
increases with
frequency.
The Visible Spectrum
 EM waves that you can see are called
visible light;
 Color
varies depending of the wavelength;
from longest to shortest.
Red / orange / yellow / green / blue /
indigo/ violet.
Explaining the Atomic Spectra
 Lowest possible energy of the electron is its
ground state.
 Energy
raises (excites) electrons from ground state
and is also released (light) when they drop back
(quanta of light released is called a Photon.)
 Bohr Model of Hydrogen (ref. table)
The light emitted by an electron moving from a
higher to a lower energy level has a frequency
directly proportional to the energy change of the
electron.