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Chapter 4
Electrons
In
Atoms
Chapter 4
Section 1
New Atomic
Model
Objectives
Explain the mathematical relationship
among speed, wavelength and frequency
of electromagnetic radiation.
Discuss the dual wave-particle nature of
light.
Describe the photoelectric effect.
Describe the Bohr model of the atom.
Rutherford Model
Was an improvement over previous models.
Helped to explain the positively charged
nucleus.
It did not explain where the atom’s negatively
charged electrons are located in space
around the nucleus.
Light and Electrons
To begin to grasp the nature of electrons,
examining the nature of light is necessary.
We will begin by first introducing some
properties of light.
We will then see how these properties are
related to the properties of the electron.
Properties of Light
Light behaves as waves and has wave-like
properties.
Electromagnetic Radiation – a form of
energy that exhibits wavelike behavior as it
travels through space.
Kinds of electromagnetic radiation include
visible light, X rays, ultraviolet and infrared
light, microwaves and radio waves.
Properties of Light
Electromagnetic Spectrum – includes all forms
of electromagnetic radiation.
Electromagnetic Spectrum
Properties of Light
All forms of electromagnetic radiation move
at a constant speed.
3.0 x 108 m/s
This is considered the speed of light.
A significant feature of waves is its repetitive
nature.
Waves can be characterized by two features:
Wavelength (l) - the distance between
corresponding points on adjacent waves.
The units for wavelength are meters,
centimeter and nanometer depending on the
form of electromagnetic radiation.
Wavelength
Frequency (n) – defined as the number of
waves that pass a given point in a specific
time, usually one second (hertz - Hz).
Frequency and wavelength are related by
the following equation:
c = ln
c = speed of light
l = wavelength
n = frequency
c = ln
Because c is the same for all
electromagnetic radiation, the product ln
is a constant.
l is inversely proportional to n
As the wavelength (l) of light increase, its
frequency (n) decreases, and vice versa.
The Photoelectric Effect
Photoelectric Effect – refers to the emission
of electrons from a metal when light shines
on the metal.
When light strikes a metal, no electrons
were emitted if the light’s frequency was
below a certain minimum.
Photoelectric Effect Experiment
The Photoelectric Effect
The explanation for the photoelectric effect is
attributed to German physicist Max Planck.
Planck proposed the energy needed to eject
an electron from the surface of an object is:
E = hn
The Photoelectric Effect
E = hn
E = energy
n = frequency
h = Planck’s constant – 6.626 x 10-34 J-s
The Photoelectric Effect
This energy can also be related to its
wavelength by the following equations:
E = hn and c = ln
to get:
hc
E=
l
The Photoelectric Effect
Albert Einstein expanded on Planck’s theory
by explaining that electromagnetic radiation
has a dual wave-particle nature.
Light can also be thought of as a stream of
particles.
Each particle of light carries a quantum of
energy.
The Photoelectric Effect
Einstein called these particles photons.
Photon – a particle of electromagnetic
radiation having zero mass and carrying a
quantum of energy.
The energy of a particular photon depends
on the frequency of radiation:
Ephoton = hn
The Photoelectric Effect
Summary:
Light has both wave properties (l and n)
and particle (photons) properties.
In order for an electron to be ejected from a
metal surface, the electron must be struck
by a single photon possessing the minimum
energy (frequency and wavelength).
Atom Line Emission
Spectrum
Ground State – the lowest energy state of
an atom.
Excited State – A state in which an atom
has a higher energy than it has in its
ground state.
When an excited atom returns to its ground
state, it gives off energy that it gained in the
form of electromagnetic radiation.
E2
Electric current
E1
Excited state energy
Electromagnetic radiation
Ground state energy
When an electric current was passed through
a tube containing hydrogen gas, a pink glow
of light was emitted.
When this pink emitted light was passed
through a prism, it was separated into a
series of specific wavelengths of visible light.
The bands of light were part of what is known
as hydrogen’s line-emission spectrum.
Hydrogen Atom Line Emission
Spectrum
Why has the hydrogen atoms given off only
specific wavelengths of light?
Scientists had expected to observe the
emission of a continuous range of
wavelengths of electromagnetic radiation,
that is a continuous spectrum.
Whenever an excited hydrogen atom falls
back from an excited state to its ground state,
it emits energy (E).
The energy is: E = hn
This energy is equal to the difference in
energy between the atom’s excited state (E2)
and its ground state (E1).
E2 – E1 = Ephoton = hn
Energy difference between
ground and excited state
The fact that hydrogen atoms emit only
specific wavelengths of light indicated that
the energy differences between the atom’s
energy states were fixed.
This suggested that the electron of a
hydrogen atom exists only in very specific
energy states.
Bohr Model of the Hydrogen Atom
Niels Bohr, a Danish physicist explained the
line spectrum of hydrogen in 1913.
His model combined the concepts of Planck
and Einstein. E = hn
Bohr assumed the atom contained a nucleus
and that the electrons circled the nucleus in
circular orbits.
Bohr Model
The three postulates of the Bohr model:
1) The electron in the hydrogen atom may
only occupy orbits of certain radii that
correspond to certain discrete energies.
2) While an electron is in an allowed energy
orbit, it does not radiate energy and it
remains in that orbit without crashing into
the nucleus.
Bohr Model
3) An electron may move from one energy
state to another by absorbing or releasing
energy. The energy needed is the
difference between one energy level and
another and is equal to a photon,
Ephoton = hn
Bohr Model
release energy
ground state
+
eabsorb energy
excited state
Bohr Model
Bohr Model
Ephoton = hn
By knowing the wavelengths from the
hydrogen atom line emission spectrum,
Bohr could solve for the energy of the
photon using the above equation.
This energy (Ephoton) represents the
difference in energy between the different
orbits of the hydrogen atom.
Bohr Model
While the Bohr model works well for
hydrogen, it does have its limitations:
1) It did not work well with atoms with more
than one electron.
2) It does not account for electron-electron
repulsions.
3) Additional electron-nucleus interactions
present problems.
Section Review
1. What is the major shortcoming of the
Rutherford model of the atom
2.Write an equation that relates speed of
light, wavelength and frequency
3. Write equations that relate energy to either
frequency of light or wavelength of light
Lab Demo
light experiments with
various gases
Chapter 4
Section 3
Electron
Configurations
Objectives
List the atomic orbitals of an atom.
List the total number of electrons needed to
fully occupy each main energy level.
State the Aufbau principle, the Pauli
Exclusion principle and Hund’s rule.
Write the electron configuration for any
element.
Atomic Orbitals
Wave Model
A more complex, highly mathematical model
was developed to explain observations of
atoms containing more than one electron.
This model works for all the elements and not
just for hydrogen as in the Bohr model.
Electronic Configuration – describes the
arrangement of electrons in an atom.
Because atoms of different elements have
different number of electrons, a distinct
electron configuration exists for each
element.
The electrons will assume arrangements
that have the lowest possible energies.
Ground State Configuration – the lowest
energy arrangement of the electrons for
each element.
Atomic Orbitals
Bohr Model – the orbit of the electron was
circular around the nucleus.
In the wave model the simple circular orbit
was replaced with 3D orbitals (electron
clouds) of various shapes in which an
electron is likely to be.
Atomic Orbitals
There are four main atomic orbitals which
describe the electron configuration of the
elements:
S orbital - spherical shape
P orbital – dumbbell shape
D orbital – clover shape
F orbital – Too complex to discuss.
Periodic Table with Orbitals
S orbital - spherical shape
P orbital – dumbbell shape
D orbital – clover shape
s, p and d orbitals
Atomic Orbitals
Energy levels of the three orbitals of interest:
S orbital – lowest energy
P orbital – slightly higher in energy
D orbital – higher in energy than P orbital
Electron Configuration Rules
The number of electrons in an atom is the
same as the number of protons.
So the periodic table will be of real value in
determining electron configurations.
To build up electron configurations for any
particular atom, first energy levels of the
orbitals are determined.
Energy
4p
3d
4s
3p
3s
2p
2s
1s
Nucleus
Electron Configuration Rules
The electrons are added to the orbitals one
by one according to three basic rules:
1) Aufbau Principle – An electron occupies
the lowest energy orbital that can receive
it.
The orbital with the lowest energy is the 1s
orbital. The one electron of hydrogen goes in
this orbital.
Electron Configuration Rules
The 2s orbital is the next highest in energy,
then the 2p orbitals.
The numbers 1,2,3 etc. refer to the row of
the periodic table the atom is located in.
As can be seen on the diagram there is
only 1-s orbital, 3-p orbitals and 5-d orbitals.
These refer to their orientation in space.
Electron Configuration Rules
Note on the energy level diagram that the 4s
orbital is lower in energy than the 3d orbital.
Therefore, the 4s orbital is filled before any
electrons enter the 3d orbitals.
Aufbau Principle
4d
4s
3s
2s
1s
4p
3p
2p
3d
Electron Configuration Rules
2) Pauli Exclusion Principle – no more than
two electrons may be present in an orbital and
their spins must be paired.
This rule basically states no two atoms can have
the same electron configurations.
Electron Configuration Rules
3) Hund’s Rule – orbitals of equal energy are
each occupied by one electron before any
orbital is occupied by a second electron. The
spins of these electrons must be opposite.
This rule is because similarly charged
electrons want to be as far away as possible.
2p orbital
Energy Level Diagram for Oxygen
3p
energy
3s
2p
2s
1s
nucleus
Energy level diagram for oxygen
Electronic Configurations
Electronic configurations are important in
chemistry:
1) To predict what type of bonding will occur
with a particular element and which
electrons are being used in the bonding.
2) Helps explain the properties of elements.
Energy
4p
3d
4s
3p
3s
2p
2s
1s
Nucleus
Electronic Configurations
While energy level diagrams are very useful they
are bulky to work with.
Electron configuration notations are simpler and
give the same information.
Electronic Configurations
Electron configuration notations eliminate the
lines and arrows of the diagrams.
Instead the number of electrons in an energy
level is shown by adding a superscript to the
energy level designation.
Example: hydrogen - 1S1
Electronic Configurations
Example: hydrogen - 1S1
The large 1 indicates hydrogen is in the first
row of the periodic table. First energy level.
The S indicates the electron is in the s orbital.
The superscript 1 indicates that there is one
electron in the 1S orbital.
Electronic Configurations
Example: helium - 1S2
The superscript 2 indicates that there are two
electrons in the 1S orbital.
Problem: Give the electron configuration of
boron and explain how the electrons are
arranged.
Elements of the Second Period
In the first period elements, hydrogen and
helium, electrons occupy the first energy level –
1s.
After the 1s orbital is filled, the next electron
occupies the 2s orbital – Aufbau principle.
Lithium has an electron configuration of 1s22s1
Aufbau Principle
4d
4s
3s
2s
1s
4p
3p
2p
3d
Classwork
Page 107 - Problems 1 – 2
Homework
Energy Level Diagram Worksheet
Elements of the Second Period
Highest Occupied Level – is the electron
containing main energy level with the largest
number.
In the case of lithium that is the 2s level.
Inner Shell Electrons – The electrons which are
in the levels below the highest occupied level.
In the case of lithium that is the 1s level.
Elements of the Second Period
Elements of the Second Period
When you get to neon (Ne) all the 2s and 2p
orbitals are full.
Octet Rule – when all of the sublevels (s and
p orbitals) of the highest occupied level is
filled with eight electrons.
All the elements in the last column of the
periodic table obey the octet rule.
Noble Gases
Neon is a member of the Group 18 elements
(last column).
These elements include neon, argon, krypton,
xenon and radon).
These elements are known as the noble gases.
Elements of the Third Period
Noble Gas Configuration
To simplify sodium’s notation, the symbol for
neon, enclosed in brackets, is used to represent
the complete neon configuration.
[Ne] = 1s22s22p6
So the electron configuration for sodium can be
written:
[Ne]3s1
This is the noble gas configuration
Elements of the Fourth Period
With the 4s level full (calcium), the 4p and 3d
sublevels are next available.
Referring to the Aufbau diagram of energy
levels, the 3d sublevel is lower in energy than
the 4p sublevel.
There are five 3d orbitals that hold a total of 10
electrons. Elements range from Sc to Zn.
Elements of the Fourth Period
Elements of the Fifth Period
Elements in the fifth period start with the 5s
orbital.
5s
4d
5p
Periodic Table with Orbitals
Problem
1) Write the energy level configuration, the
electron configuration and noble gas
configuration for iron (Fe).
2) How many orbitals are filled? How many
unpaired electrons are there in an atom of
iron?
Classwork
Page 115
Practice Problems 1 – 3
Page 116
Practice Problem 1
Homework
Worksheet
Electron Configuration/Noble Gas
Configuration