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Chapter 5
Structures of the Atom
Electron
Electrons were discovered in 1897 by Sir J. J. Thomas
Electrons have the following characteristics:
Travel from a cathode (electron emitter) in straight lines.
Cause a piece of metal foil to become hot by striking it for a period of
time.
Deflected by magnetic and electric fields showing that they have a
negative charge.
Are emitted by some metals when heated. This is called “thermal
emission”.
Are emitted by some metals when exposed to light. This is called
“photoelectric effect”.
Proton
Protons discovered in 1886 by Eugene Goldstein using a Crookes Tube
Protons have the following characteristics:
Positive charge
same strength as an electron
found in nucleus
Neutron
Neutrons were discovered in 1932 by James Chadwick
Neutrons have the following characteristics:
mass almost identical to proton
neutral charge
will disintegrate to form a proton and an electron when outside the
nucleus
Radioactivity and the Atomic Structure
In 1896 Antoine Becquerel discovered that certain elements decomposed
to form other elements by emitting radiation.
There are three types of radiation
alpha () particle which is a He+2 ion
beta () particle which is an egamma () ray which is made up of high energy and has no charge
Nuclear Atom
In 1911 John Rutherford proposed that the atom was composed of a very
small, positively charged nucleus where most of the mass is concentrated,
surrounded by the number of electrons necessary to produce an
electrically neutral atom.
He was able to propose this by bombarding a thin gold foil with
alpha particles.
Because only a few alpha particles were deflected, Rutherford
proposed that there must be a small nucleus within the atom.
This is today’s model of the atom.
Bohr Model of the Atom
If electrons spin around the nucleus, why don’t they eventually slow
down and get sucked into the nucleus?
This was the question posed by many scientists.
Niels Bohr proposed that the electrons are quantized, or restricted
to discrete, individual areas or values.
The energy of the electron is restricted to certain values which
correspond to the radius of its orbit.
The radius is characterized by an integer, n.
The equation being:
E   kZ2
2
n
where k = 2.179x10-18J; Z = the atomic number; and n = the integer characteristic
of the orbit.
Atomic Spectra and Atomic Structure
An electron in an atom can be given enough energy to be promoted from
one orbital to the next higher.
When that electron moves back to its original orbital, that added energy is
given off in the form of light.
The amount of energy given off can be determined in the following
equation where E is energy in Joules.
 1
1 
E  2.179 x1018  2  2 
 n1 n2 
n1 is the lower-energy inner orbital
n2 is the higher-energy outer orbital
The wavelength of the emitted energy can be calculated by using
the formula:
 25

hc
1.98647x10
or  
E
E
The Quantum-Mechanical Model of the Atom
Erwin Schrödinger showed that the electron may be visualized as being in
rapid motion within one of several regions of space located around the
nucleus.
He called these regions “orbitals”
Some scientists call the occupancy of these orbitals “electron densities”
because of the rapidity of the electron moving within.
The location of an electron cannot be determined exactly.
We can identify only the region or volume of space where there is a
relatively high probability of finding the electron
Results of the Quantum-Mechanical Model of
the Atom
Orbitals are characterized by n, the principal quantum number, which
may take on any integral value: n = 1, 2, etc.
As n increases, the orbitals extend further away from the nucleus.
A shell contains orbitals with the same value of n.
A shell will have the values of K, L, M, etc.
Shell
K
L
M
N
n
1
2
3
4
Orbitals with the same value of n may have different shapes called the
azimuthal quantum number distinguished by the second quantum
number l.
An electron with l = 0 will have a spherical orbital shape or s-orbital
An electron with  = 1 will be a p-orbital
An electron with  = 2 will be a d-orbital
An electron with  = 3 will be an f-orbital
The shapes of the f-orbitals can be found in your text.
So:
Shell
K
L
M
N
O
n Value
1
2
3
4
5
Possible  Values
0
0, 1
0, 1, 2
0, 1, 2, 3
0, 1, 2, 3, 4
Types of Orbibals
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f 5g
For  values larger than zero, orbitals with the same value of  in any
given shell have the same shape and the same energy but differ in their
orientation.
There are 2 + 1 orbitals with the same  quantum number that differ in
their orientation
Each orientation is characterized by a third quantum number, m, the
magnetic quantum number.
When  = zero, there is only one orientation in space for the orbital
When  = 1, there are three possible values of m (-1, 0, +1) and thus the porbital has three sub-orbitals.
When  = 2, there are five possible values of m (-2, -1, 0, +1, +2) and thus
the d-orbital has five sub-orbitals.
Electrons have a fourth quantum number, s, called spin quantum number.
The spin specifies the direction of spin of an electron about its own axis.
It can be either clockwise or counterclockwise designated as s= +1/2 or s=
-1/2
The energy of an electron in an atom is limited to discrete values.
The energies can be determined from the wave function that describes the
behavior of the electron.
The maximum number of electrons that may be found in a shell with a
principal quantum number of n is 2n2.
The maximum number of atomic orbitals in a shell is given by n2.
Orbital Energies and Atomic Structure
The energies of atomic orbitals increase as the principal quantum number,
n, increases.
The energies of the orbitals increase within a shell as the quantum
number, , increases.
For atomic numbers greater than 20, the relative energies of the orbitals
may differ slightly from the order shown.
For example, the energies of the electrons in 1s orbitals become lower and
lower as the atomic numbers of the atoms increase.
It is possible to predict the electron configuration of an atom in its ground
state using the following guidelines:
Electrons occupy the lowest possible energy levels or orbitals
The Pauli Exclusion Principle states that no two electrons in the
same atom can have the same set of four quantum numbers.
Hund’s Rule: Every orbital in a subshell is singly occupied with
one electron before any one orbital is doubly occupied, and all
electrons in singly occupied orbitals have the same spin.
The Aufbau Process
The systematic variations in the electronic structures of the various
elements can be “built” in atomic order.
This process is called the Aufbau process
Note the exceptions to the aufbau process found in your text book.
Some of the exceptions are chromium, copper and lanthanum.
These exceptions are usually due to very similar energies within
the subshells.