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Chapter 12
Electrons in Atoms
Introduction

The view of the atom as a positively charged
nucleus (protons and neutrons) surrounded by
electrons is useful for visualizing the basic
structure of an atom
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However, this model of the atom explains only a few
simple properties of atoms.
This chapter will include models of atomic
structure with an emphasis on the electrons in
atoms
12.1 The Development of Atomic Models

For about 50 years after the time of Dalton
(1766-1844) the atom was considered to be a
solid indestructible mass

The discovery of subatomic particles shattered every
theory that included indestructible atoms
12.1 Thompson Model

After discovering the
electron, Thomson
described a “plum
pudding” model of the
atom in which an atom
was a ball of positive
charge containing
electrons.
12.1 Rutherford Model

After learning of the
nucleus, Rutherford
proposed the nuclear
atom in which electrons
surround a dense nucleus
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most of the atom being
empty space.
Later experiments showed
that the nucleus is
composed of protons and
neutrons
12.1 Bohr Model

In 1913, Niels Bohr
proposed the “planetary
model” of the atom in
which the electrons
move in orbits around
the nucleus.
12.1 Bohr Model
Bohr proposed that
electrons in a particular
path have a fixed energy
that keeps them from
falling into the nucleus.
energy level – the region
around the nucleus where
the electron is likely to be
moving
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Ladder Analogy
12.1 Bohr Model

Electrons can jump from one energy to another
by gaining or losing just the right amount of
energy
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Electrons cannot exist between energy levels
A quantum of energy is the amount of energy required
to move an electron to the next highest energy level.
The higher the energy level the farther the electron is from
the nucleus (usually)
Energy levels are more closely spaced further from the
nucleus
The higher the energy level the easier it is for the electron
to escape
12.2 Quantum Mechanical Model

Quantum mechanical
model states that the
atom has no definite
shape and that electrons
do not have precise
orbits
12.2 Quantum Mechanical Model

In 1926, Erwin Schrödinger used the new
quantum theory to write and solve a
mathematical equation to describe the location
and energy of an electron in a hydrogen atom

The modern description for electrons in the
atom, the quantum mechanical model,
comes from the mathematical solution to
Schrödinger’s equation.

Primarily mathematical – has few analogies in the
visible world
12.2 Quantum Mechanical Model

Features of the quantum mechanical model:
restricts the energy of electrons to certain values but
does not define the exact path taken by the electron
 estimates the probability of finding the electron
within a given region called a cloud (electron cloud)
 the electron can be found within this cloud 90% of
the time
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12.3 Atomic Orbitals
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The quantum mechanical model energy levels
with principle quantum numbers (n)
Each principle quantum number refers to a
major (principle) energy level in an atom
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n=1, 2, 3, 4, etc.
The average distance of the electron from the
nucleus increases with increasing values of n.
12.3 Atomic Orbitals

Within each principle energy level, the electrons
occupy energy sublevels

The number of energy sublevels is the same as the
principle quantum number
12.3 Atomic Orbitals

The quantum mechanical model describes the
position of an electron with cloud shapes (based
on probability)
These cloud shapes are called atomic orbitals
 Atomic orbitals are represented with letters: s, p, d, f, g
 Regions where there is a low probability to find
electrons are called nodes
 Atomic orbitals have different characteristic shapes

s orbitals are spherical
 p orbitals are dumbell shaped
 d and f orbitals are more complex and harder to visualize
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12.3 Atomic Orbitals
12.3 Atomic Orbitals
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Table 12.1 Summary of Principle Energy Levels, Sublevels,
and Orbitals
Principle Energy
Level
# of
Sublevels
Type of Sublevel
n =1
1
1s (1 orbital)
n=2
2
2s (1 orbital), 2p (3 orbitals)
n=3
3
3s (1 orbital), 3p (3 orbitals), 3d (5 orbitals)
n=4
4
4s (1 orbital), 4p (3 orbitals), 4d (5
orbitals), 4f (7orbitals)
12.3 Atomic Orbitals
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Each orbital can hold a maximum of 2
electrons
The maximum number of electrons that can
occupy a principle energy level is given by the
formula 2n2 where n = principle quantum
number
Example: How many electrons can be found in
the 3rd principle energy level?
A: 2n2 = (2)(32) = 18 electrons
12.4 Electron Configurations
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In all natural phenomena, change proceeds
toward the lowest possible energy state.
High-energy systems are unstable
 Unstable systems lose energy to become more stable
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In the atom, electrons and the nucleus
interact to make the most stable arrangement
possible.
12.4 Electron Configurations
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The ways in which electrons are arranged
around the nuclei of atoms are called
electron configurations.
Three rules or principles are used to determine
the electron configuration of atoms:
1.
2.
3.
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule
12.4 Aufbau Principle
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Electrons enter orbitals of lowest energy first
Orbitals within a sublevel of a principle energy level
are always of equal energy
 See Figure 12.7, page 330
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12.4 Aufbau Diagram
12.4 Pauli Exclusion Princple
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An atomic orbital may describe at most two
electrons
To occupy the same orbital, two electrons must have
opposite spins (the electrons spin must be paired)
or they would repel each other
 Electron spin can clockwise or counterclockwise and
is represented by vertical arrows that point up or
down
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Hund’s Rule
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When electrons occupy orbitals of equal energy,
one electron enters each orbital until all the
orbitals contain on electron with parallel spins
(arrows in the same direction).
Second electrons then add to each orbital so that
their spins are paired with those of the first
electrons in the orbital.
12.4 Electron Configurations

You can use a shorthand method for showing
the electron configuration of an atom
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This involves writing the energy level and symbol for
every sublevel occupied by an electron
A superscript indicates the number of electrons
occupying that sublevel.
 See Table 12.2 for examples
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12.4 Example 1
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Use Figure 12.7 to write electron configurations
and orbital diagrams for these atoms.
a. phosphorus
b. nickel
12.4 Practice Problems
8. Arrange the following sublevels in order of
decreasing energy: 2p, 4s, 3s, 3d, and 3p.
9. Write electron configurations for atoms of the
following elements. How many unpaired
electrons do these atoms have?
a. boron
b. fluorine
12.5 Exceptional Electron Config.
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Like most rules, there are exceptions to the rules
for determining electron configurations.
Cr, Cu, Ag, Au, and Pt are five of the 14
elements that fill their electron orbitals a little
differently than all the other elements
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These particular elements are more stable with an
electron moving into 3d and making 4s half filled
rather than a filled 4s shell, this is due to their
geometric structure
12.5 Exceptional Electron Config.
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Following the Aufbau diagram, this would be
the electron configuration for Cr and Cu:
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This is the correct electron configuration: