Chapter 5 Notes
Section 1 – Models of the Atom
Atomic Models
 Rutherford’s model showed that electrons
moved around the nucleus, but could not
explain the chemical properties of
elements.
 In 1913 Niels Bohr proposed that an
electron is found only in specific circular
paths, or orbits, around the nucleus
 Each electron orbit has a fixed energy,
called energy levels
 The closer the orbit to the nucleus the
lower its energy level
Atomic Models
 An electron can move from one level to
another if it gains or loses a specific
amount of energy
 A quantum is the amount of energy an
electron needs to move levels.
 The amount of energy needed to move
levels isn’t always the same.
 The higher levels are closer together so it
takes less energy to move between those
levels
The Quantum Mechanical Model
In 1926, Erwin Schrodinger devised
and solved a mathematical equation
describing the behavior of the
electron in a hydrogen atom
Resulted in the Quantum Mechanical
Model
This model determines the allowed
energies an electron can have and how
likely it is to find the electron in various
locations around the nucleus.
Atomic Orbitals
 For each energy level, the Schrodinger
equation also leads to a mathematical
expression called an atomic orbital
 An atomic orbital is the region of space in
which there is a high probability of finding
an electron
 Each energy level is labeled by principal
quantum numbers (n).
 N = 1, 2, 3, …
 Each energy level can have several
orbitals
Atomic Orbitals
Each energy sublevel corresponds to
an orbital of a different shape,
which describes where the electron
is likely to be found
s orbitals are spherical
p orbitals are dumbbell-shaped
d and f orbitals are slightly more
complicated in shape
Each orbital can hold a maximum of
2 electrons
Atomic Orbitals
 The lowest principal energy level (1) has 1 sublevel – 1s
 The second principal energy level (2) has 2 sublevels – 2s,
2p
 The third principal energy level (3) has 3 sublevels – 3s,
3p, 3d
 The fourth principal energy level (4) has 4 sublevels – 4s,
4p, 4d, 4f
 The # of orbitals varies with each sublevel
 s – 1, p – 3, d – 5, f – 7
 To find the maximum number of electrons use the formula
2n2 (n is the quantum energy level)
Chapter 5 Notes
Section 2 – Electron Arrangement in Atoms
Electron Configurations
The ways in which electrons are
arranged in various orbitals around
the nuclei of atoms are called
electron configurations
There are three rules on how to find
the electron configuration of an
element:
1. Aufbau principle
2. Pauli exclusion principle
3. Hund’s rule
Aufbau
Electrons occupy the orbitals of lowest
energy first.
The orbitals for any sublevel of a principal
energy level are always of equal energy
The s sublevel is always the lowest-energy
The range of energy levels within a
principal energy level can overlap the
energy levels of another principal level
Ex– 4s orbital is lower in energy then the 3d
orbital
Pauli Exclusion Principal
An atomic orbital may describe at
most two electrons
To occupy the same orbital, two
electrons must have opposite spins.
Spin is a quantum mechanical
property of electrons and may be
thought of as clockwise or
counterclockwise.
Vertical arrows indicate the direction
of the spin ↑↓
Hund’s Rule
 Hund’s Rule states that one electron will
enter each orbital until all the orbitals
contain one electron with the same spin
direction
 Electrons occupy orbitals of the same
energy in a way that makes the number of
electrons with the same spin direction as
large as possible
 Ex.
 If there are three electrons in the p orbital they
would fill like - ↑ ↑ ↑_
 If there are four electrons in the p orbital they
would fill like - ↑↓ ↑ ↑_
 Table 5.3
Exceptional Electron Configurations
 Some actual electron configurations differ
from those assigned using the aufbau
principle
 Half-filled sublevels are not as stable as
filled sublevels, but they are more stable
then other configurations
 Exceptions include Copper and Chromium
and elements in the same column.