Atomic Theory: History
Dalton (1805)
Empirical knowledge
 law of definite proportions
 law of conservation of mass
Thomson (1897)
Empirical knowledge
 electrical nature of solutions
 cathode ray tubes
Rutherford (1911)
Empirical knowledge
 gold foil experiment
Chadwick (1932)
Empirical knowledge
 artificial fusion reactions
27
Al +  
30
Theoretical knowledge
Matter is composed of indestructible, indivisible
atoms, which are identical for one element, but
different from other elements.
Theoretical knowledge
Matter is composed of atoms that contain
electrons embedded in positive material. The kind
of element is characterised by the number of
electrons in the atom.
Theoretical knowledge
An atom is composed of a very tiny nucleus,
which contains positive charges and most of the
mass of the atom. Very small negative electrons
occupy most of the volume.
Theoretical knowledge
The existence of neutral particles called neutrons
in the nucleus of atoms was theorized.
P + n
The Rutherford Model of the Atom
proton (p+)
electron (e-)
neutron (n)
isotopes





a positively charged subatomic particle found in the nucleus of atoms
a negatively charged subatomic particle found in the orbitals of atoms
a neutral or uncharged subatomic particle present in the nucleus of atoms
a variety of atoms of an element; atoms of this variety have the same number of protons
as all atoms of the element but different numbers of neutrons
An atom is made up of an equal number of negatively charged electrons and positively charged
protons.
Most of the mass of the atom and all of its positive charge is contained in a tiny core region called
the nucleus.
The nucleus contains protons and neutrons that have approximately the same mass.
The number of protons is called the atomic number (Z).
The total number of protons and neutrons is called the mass number (A).
Quantum Theory
Electromagnetic spectrum includes all forms of electromagnetic radiation, from very short
wavelength gamma () rays to ordinary visible light to very long wavelength radio waves.
The photoelectric effect is the release of electrons from a substance due to light striking the surface of
a metal.

The energy is directly related to the frequency of the electromagnetic radiation, shown by the
Planck equation:
E = hf


E is energy in joules (J)
f is frequency in hertz (Hz)
h is Planck’s constant (6.6 x 10-34 J/Hz)
The greater the quantum of energy, the more likely electrons will be released from the metal.
Depending on the metal, electrons are released with varying degrees of energy.
quantum (plural, quanta) is a small discrete,
indivisible quantity; whereas a quantum of light
energy is called a photon
 photons of infrared light have small packets of
energy, or quanta
 photons of ultraviolet light have large quanta
Atomic Spectra
Spectroscopy
 a technique for analyzing electromagnetic radiation (UV, IR, visible light, X-ray) spectra


bright-line spectrum produce a
series of bright lines of light emitted
by a gas, which has been excited by
heat or electricity
absorption (dark-line) spectrum
produce a series of dark lines or
missing parts in a continuous
spectrum
For a given element, only certain frequencies are absorbed (or emitted), which corresponds to
particular quanta of light.
Bohr’s First Postulate
Electrons do not radiate energy as they orbit the nucleus.
Each orbit corresponds to a state of constant energy, called a stationary (or ground) state.
Bohr’s Second Postulate
Electrons can change their energy only by undergoing a transition from one stationary state to
another.
In a transition, the quantum of energy must match the
difference in energy of the two electron states.
 these energy differences give rise to the lines in atomic
spectra
Bohr was successful at theoretically calculate the specific
wavelengths released in the atomic spectra for hydrogen.
Unfortunately for more complex elements, Bohr could not
predict the wavelength patterns that were empirically
identified.
Quantum Numbers
Principal Quantum Number, n
 relates primarily to the main energy that an electron possesses;
n = 1, 2, 3, 4…
 the main energy level of electrons (the circles drawn for a Bohr diagram)
Secondary Quantum Number, l
 relates primarily to the shape of the electron orbit
 corresponds to a sublevel within the main energy level
 the number of values for l equals the volume of the principal quantum number (in other words,
the number of sublevels in each main energy level arises from the value of n)
Principal
quantum
number, n
1
2
3
4

Possible secondary
quantum numbers, l
Number of sublevels
per primary level
0
0, 1
0, 1, 2
0, 1, 2, 3
1
2
3
4
each value of l relates to the shape of the electron sublevel and will help to explain the regions of
the periodic table
value of l
letter
designation
name
designation
0
1
2
3
s
p
d
f
sharp
principal
diffuse
fundamental
Each shape arises from a standing wave pattern that develops when the correct energy exists for the
electron to vibrate three-dimensionally in space and to create the characteristic shape of the sublevel.
Quantum Mechanics


Louis de Broglie (1923) first proposed the idea of an electron behaving as
a wave.
Erwin Schrodinger described more precisely the behaviour of electrons in
the atom.
The current theory of atomic structure based on wave properties of electrons
is known as quantum mechanics.
Magnetic Quantum Number, ml


relates primarily to the direction of the electron orbit
the number of values for ml is the number of
independent orientations of orbits that are possible
within each sublevel
Value of l
0
1
2
3
Values of ml
-l to +l
0
-1, 0, +1
-2, -1, 0, +1, +2
-3, -2, -1, 0, +1, +2, +3
For each sublevel (s-p-d-f), the number of possible
orientations for the orbitals is indicated by the number of
values for ml.
Value of l
0
1
2
3
Sublevel symbol
s
p
d
f
Number of orbitals
1
3
5
7
Empirical evidence such as ionization energy values and valences provide support for the theoretical
knowledge of energy sublevels and the orientations of electrons within each sublevel.
Electron orbitals are regions of space around the
nucleus where an electron is likely to be found.
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


3D region of space
distance from nucleus varies
variety of shapes
2 electrons per orbital
Heisenberg uncertainty principle states that it is impossible to simultaneously know the exact
position and speed of a particle.
A mathematical or graphical representation of the chance of finding an electron in a given space is
called electron probability density.
Spin Quantum Number, ms


relates to a property of an electron that can best be described as its spin
the value of ms can only be +1/2 or –1/2

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with opposite spins, electrons that are paired up
experience no magnetic capacity
unpaired electrons are able to exhibit magnetic properties
paramagnetism occurs when substances are weakly
attracted to magnets, while ferromagnetism describes
substances that are strongly attracted to magnets
Drawing Energy-Level Diagrams
The aufbau principle requires that each electron is
added to the lowest energy orbital available in an atom
or ion.
The Pauli exclusion principle requires that only two
electrons with opposite spins can occupy any one
orbital.
Hund’s rule requires that the orbitals in the same
energy level must have one electron before a second
electron can be placed in any orbital at that energy
level.
Rules for drawing energy-level diagrams:
1. Start adding electrons into the lowest energy level
and build up form the bottom until the limit on the
number of electrons for the particle is reached.
2. No two electrons can have the same four quantum numbers; if an electron is in the same orbital
with another electron, it must have opposite spin.
3. No two electrons can be put into the same orbital of equal energy until one electron has been put
into each of the equal-energy orbitals.
4. For anions, add extra electrons to the number for the corresponding atom. For cations, do the
neutral atom first, then subtract the required number of electrons from the orbitals with the highest
principal quantum number, n.

Electron Configuration
a method for communicating the location
and number of electrons in electron energy
levels and sublevels
Step 1
Step 2
Step 3
Find the element in the
periodic table and determine
the total number of electrons in
the atom or simple ion.
Assign electrons in increasing order of the main energy levels and sublevels using the
periodic table.
For anions, add extra electrons to the total number in the atom. For cations, use the
electron configuration for the neutral atom and remove the required number of electrons
from the highest principal quantum number, n.
A shorthand form of electron configurations is widely accepted. The core electrons are represented by
a noble gas and only the valence electrons are shown.
Electron configurations demonstrate magnetism and help explain charges for transition metals.
Explaining the Periodic Table
electron distribution
period
# of elements
1
2
3
4-5
6-7
2
8
18
18
32
Representative
elements are those
elements filling their s
or p orbital.
Transition elements
are those elements
filling their d orbital.
Lanthanide and
actinide series are
those elements filling
their f orbital.
groups:
orbitals:
1-2
s
2
2
2
2
2
13-18
p
3-12
d
-----f
6
6
6
6
10
10
10
14