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Electrons
in Atoms
Chapter 5
?
Introduction
• The atom evolved over the years as a
result of new information that was
collected.
• The quantum mechanical model of the
atom helped explain how electrons
traveled around the nucleus and their
energy.
• Light and energy are related to each other
and the movement of electrons.
• Electrons are positioned in a particular
way around the nucleus.
Models of the Atom
(Section 5.1)
• The
Development of
Atomic Models
• The Bohr Model
• The Quantum
Mechanical
Model
• Atomic Orbitals
I.) The Development of Atomic
Models
• The atom ~ 1803
• Dalton’s atom
• Small indivisible
Deficiency:
No subatomic
structures
• Represents the
smallest part of
matter that still has
unique properties
• The atom ~ 1897
• Thomson’s atom
Deficiency:
No nucleus and
no movement of
electrons
• negatively charged
electrons embedded
on a positive
charged sphere
Deficiency:
Could not explain
chemical properties of
elements such as…
• The atom ~
1911
• Rutherford’s
atom
• Small, dense,
positively
charged
nucleus with
electrons
orbiting around
it
Hydrogen
Helium
Why heated objects
glow red, then
yellow, then white.
Lithium
Why elements give characteristic colors
when heated.
II.) The Bohr Model
• A student of
Rutherford’s
• Believed
Rutherford’s model
needed to be
improved to account
for new discoveries.
• Considered the
simplest atom, the
hydrogen atom.
Niels Bohr
Danish Physicist
(1885 – 1962)
• Bohr also
proposed
electrons orbit
around nucleus
• Electrons orbit in
fixed energy
levels.
• Electrons can
move up and
down energy
levels.
• Energy is
involved in this
movement
Quantized Energy
Quantum of Energy: The amount of energy
required to move an electron from one
energy level to another.
• The Bohr atom gave
results that agreed with
experiments for the
hydrogen atom only.
How does this
energy relate to
light and color?
Light
• Dual Nature of Light: Light can act like waves,
and as straight line particles.
• Light is one type of electromagnetic radiation
(em), which is a form of energy that has
wavelike behavior
• Other types of em radiation are: x-rays, uv,
infrared, microwaves & radio, and together
they all form the Electromagnetic Spectrum
Energy and Light
The Wave
The Electromagnetic Wave
• The speed of all em waves through a
vacuum (space) and through air is:
Speed of light (c) = 3.0 x 108 m/s
• The length of each individual wave is
known as it’s wavelength () which is the
length between corresponding points on
adjacent waves, usually measured in
nanometres (1nm = 10-9m)
• The frequency (f or ) is how many waves
pass a particular point in a second and is
measured in waves/second = Hertz (Hz)
The Light Wave
Relationship Between Frequency &
Wavelength
The relationship between them is as follows:
c=fx
• What is the wavelength & color of light that
has a frequency of 6 x 1014 Hz (1/s)?
f = 6 x 1014 Hz (1/s)
c = 3 x 108 m/s
 =?
c=fx
 = c = 3 x 108 m/s = 5 x 10-7 m
f
6 x 1014 1/s
= 500 x 10-9 m
= 500 nm ( = green )
Energy, Wavelength, &
Frequency
• When an object gets hot, it emits energy in
small specific amounts called quanta.
• A quantum is the minimum amount of energy
that can be lost or gained by an atom.
• The frequency determines the Energy…
The frequency determines the Energy by:
E = h.f
or
E = h (c/λ)
h = 6.626 x 10-34Js
(Planck’s constant)
What is the Energy of Green light with a
frequency of 6 x 1014 Hz ?
f= 6 x 1014 Hz (1/s)
E = h.f
= 6.626 x 10-34 J.s (6 x 1014 Hz)
= 3.98 x 10-19 J
The Particle Nature of Light
•
•
•
•
It was Einstein that proposed that light
acted like a stream of particles called
photons.
A photon is a particle of em radiation that has
zero mass and carrying a quantum of energy
There must be a minimum amount of energy to
eject electrons from a metal
Using E = h.f, Planck realized that the em
radiation providing the energy, must be of a
certain frequency.
Different metals need different minimum
amounts of energy, and therefore frequencies
• When the electron falls
back to its ground state
or a lower energy state,
it emits a photon of
radiation with a specific
amount of energy, and
not continuous
amounts of energy.
• Therefore, specific
frequencies of light are
emitted when excited
hydrogen “cools” down.
• When the H atom gets
excited, the e- jumps to
another specific orbit and
not somewhere in between
– like going up a ladder
• When the e- falls back, it
loses energy in the form of
a photon (a bundle of light
energy)
Convert circular
orbits to lines
Energy Levels of an Atom
The Absorption & Emission Spectrum
Emission Spectrum: The pattern of discrete
lines resulting from the frequencies of light
emitted by an element that has been energized.
• The emission
spectra
represents the
movement of
an electron
from higher
energy levels
down to lower
energy levels.
• This spectrum
is for hydrogen.
Atomic Spectra
Hydrogen
Helium
Mercury
Argon
Iodine
Neon
Explanation of the Atomic
Spectra
• Each discrete line in an emission spectrum
corresponds to one exact frequency of light
emitted by the atom.
• Emission spectrum, like a person’s fingerprint,
can be used to ID an element.
• As we saw earlier, the light emitted by an
electron moving from a higher to a lower energy
level has a frequency and wavelength directly
proportional to the energy change of the
electron.
The Bohr/ Rutherford Planetary
Orbital Model
• Based on describing
paths of moving
electrons as you would
describe the path of large
moving objects.
What was to be
done?
• Data collected and
theoretical calculations
showed that electrons
did not move in this way
for larger atoms.
III.) The Quantum Mechanical
Model of the Atom
• Comes from
mathematical
solution to a complex
equation.
• Based on
probabilities.
• Electrons occupy
certain spaces
around nucleus.
• Used theoretical calculations and “new”
experimental results to devise a
mathematical equation describing
the behavior of the electron in
a hydrogen atom.
• The quantum mechanical model of the
atom comes from the mathematical
solution of this equation.
• Solution restricts the energy of electrons
to certain energy levels.
Erwin Schrodinger
Austrian Physicist
1887 - 1961
• Solution does not involve an exact path
the electrons travel around the nucleus.
Summarizing the Findings of
Quantum Mechanics
• Determines the allowed energies an
electron can have & how likely it is to find
the electron in various locations around
the nucleus.
• The probability of finding an electron within
a certain volume of space surrounding the
nucleus can be represented by a fuzzy
cloud.
The cloud is more dense where the probability of finding an
electron is high.
The cloud is less dense where the probability of finding an
electron is low.
Summary of Atomic Orbitals
• Defined as a region of space in which
there is a high probability of finding an
electron.
• Each energy level may possess several
orbitals with different shapes and at
different energy levels.
• These energy levels are deemed
“sublevels” and corresponds to an orbital.
The Atomic Orbitals
The s Orbital
(1shape)
Each energy sublevel
corresponds to an orbital
of a different shape.
The p
Orbitals
(3 shapes)
The d orbitals (5 shapes)
Question is: Why can’t e-’s be in an
orbit between the specific Energy
levels?
French scientist Louis DeBroglie pointed
out that the electron orbits acted like the
behavior of waves.
(i.e. you can only have a certain amount of
waves in a given distance/space – not ½
of a wave)
If you can only have specific # of waves,
then based on the formula c = f x  you
can only have specific f’s, which in turn
translates into specific Energy levels.
Each energy level has energy sublevels that
correspond to an atomic
orbital.
Principle Energy
Level
Number of
Sublevels
Type of Sublevel
n=1
1
1s (1 orbital)
N=2
2
2s (1 orbital)
2p (3 orbitals)
N=3
3
N=4
4
3s (1 orbital)
3p (3 orbitals)
3d (5 orbitals)
4s (1 orbital)
4p (3 orbitals)
4d (5 orbitals)
4f (7 orbitals)
3d
3p
3s
n=3
2p
n=2
2s
n=1
1s
Describing Energy Sublevels
• The number of
sublevels per
energy level
corresponds
to the energy
level number.
• There is much
overlap of energy
sublevels.
Fill in the Following Chart
Principle
Energy Level
n=1
n=2
n=3
Number of
Sublevels
Type of
Sublevel
• There is a set number
of orbitals per energy
sublevel that
corresponds to the
number of shapes
each orbital
possesses.
•s=1
•p=3
•d=5
•f= 7
• Only 2 e- per orbital
max
Electron Configuration in Atoms
(Section 5.2)
• Electron
Configurations
• Exceptional
Electron
Configurations
I.) Electron Configuration
• The way in which the electrons of an
atom are arranged in various orbitals
around the nuclei of atoms.
• 3 basic rules allow us to determine this:
the aufbau principle, the Pauli exclusion
principle, & Hund’s rule.
• Use these rules to place all the electrons
of an atom into the different orbitals.
The 3 Rules for Electron Configurations
The Aufbau Principle
• Electrons occupy the orbitals of lowest
energy first.
Pauli Exclusion Principle
• The maximum number of electrons per orbital is two.
• When two electrons share an orbital they must have the
opposite spins (i.e. electron spins must be paired).
Hund’s Rule
• 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.
The Aufbau Principle
Electrons
occupy the
orbitals of
lowest
energy first.
The filling of atomic orbitals
is not simple beyond the 2nd
energy level.
Pauli Exclusion Principle
• The maximum number of electrons per
orbital is two.
• When two electrons share an orbital
they must have the opposite spins (i.e.
electron spins must be paired).
↑↓
Hund’s Rule
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.
Electrons repel and do not like being around
each other, and so if they can they will separate.
2 p orbitals
The Maximum # of Electrons
That Can Occupy a Principle
Energy Level
#
e
=
2
2n
– n = the principle energy level
Let’s Start
Write the
electron
configuration for
phosphorus.
(Atomic # = 15)
How many total
electrons?
Where do the
electrons
go first?
Where do the
others go?
Let’s Practice
Write the
electron
configuration for
carbon.
(Atomic # = 6)
Write the
electron
configuration
for argon.
How many total
electrons?
Write the
electron
configuration
for nickel.
How many total
electrons?
Shorthand Method for Electron
Configuration
This method involves writing the energy level
and the symbol for every sublevel occupied by
an electron.
• We indicate the number of electrons occupying a
particular sublevel with a superscript.
• When writing shorthand, keep sublevels of a
given energy level together, even though this
distorts what is obtained by following the aufbau
principle.
Let’s try this with our previous examples.
Write the electron configuration for the
following atoms using the shorthand
method.
• Phosphorus
• Carbon
• Argon
• Nickel
Simple Table to Help Write Electron
Configurations Shorthand
Use the
table to
write the
electron
configuration
for potassium.
Let’s Practice
1.
2.
3.
4.
5.
Write the electron configuration for the
following elements using the shorthand
method.
magnesium
nitrogen
chlorine
silicon
calcium
Identifying Elements from
Electron Configurations
1. Count the total number of electrons.
2. This is the atomic number.
3. Use this number to ID element from the
periodic table.
4. Example:
1s22s22p63s1
Let’s Practice
Give the symbol and name of the
elements that correspond to the following
electron configuration.
1.) 1s22s22p63s1
2.) 1s22s22p3
Electrons
in Atoms
Chapter 5
The End
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