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Electron Configuration
Orbitals
Unit 2
Photoelectric Effect
The emission of electrons from a metal when light shines on the
metal.
Particle Description of Light
A quantum of energy is the minimum quantity of energy that can
be lost or gained by an atom.
Max Planck explained relationship between a quantum of energy
and the frequency of radiation: E=hv
E=energy in joules
v= frequency is s-1
h= Planck’s constant 6.626 X 10-34 J s
Properties of Light
 Wavelength (λ, lambda) – the distance between
corresponding points on adjacent waves.
 Frequency (v)- is defined as the number of waves that pass a
given point in a specific time, usually one second.
 These are mathematically related with c=λv
 This is the speed of light equation. c is the speed of light
(m/s) λ is the wavelength of the electromagnetic wave (in
m), and v is the frequency of the electromagnetic wave (in s1).
Wavelength and Frequency
When we refer to electrons as:
 Ground state: the lowest energy state of an atom.
 Excited state: an atom has a higher potential energy than
it has in its ground state.
 Quantum theory: the mathematical wave properties of
electrons and other very small particles.
 Emission: when electrons fall to a lower energy level, a
photon is emitted.
 Absorption: when electrons move from a lower to a
higher energy level. (Energy has to be added)
 A photon is a particle of electromagnetic
radiation with zero mass and carries a quantum of
energy.
Absorption & Emission of Photons
Bohr Model
 Niels Bohr said electrons can circle the nucleus in allowed
paths (orbits)
 The energy of electrons is higher the further away from the
nucleus.
What is your address?
 How many of you have the same..
 Zip code?
 Live on the same street?
 Have the same house number?
 Just like you can’t have the exact same address as someone an
electron can not have the same set of four quantum numbers
as another electron.
 Diffraction- bending of a wave as it passes the
edge of an object or through a small opening
 Electron beams can interfere with each other.
When waves overlap it is referred to as
Interference.
Heisenberg Uncertainty Principle
 Werner Heisenberg, German physicist, proposed
any attempt to locate a specific electron with a
photon will knock the electron off its course.
 Heisenberg Uncertainty Principle- it is
impossible to determine simultaneously the
position and velocity of an electron or any other
particle.
Orbitals
 Erwin Schrodinger, Austrian physicist (1926), stated an
equation where electrons move in waves.
 Heisenberg and Schrodinger’s theories combined laid
the foundation for the Quantum theory.
 Orbitals are three-dimensional regions around the
nucleus that provides a probability of a the location of an
electron.
 Quantum numbers- specify properties of atomic orbitals
and electrons in orbitals.
Quantum Numbers
 Principle quantum number, n, the main energy level
occupied by the electron.
 As n increases the electron’s energy and its average distance
from the nucleus increases. The total number of orbitals that
exist in a given shell, or main energy level, is equal to n2.
 Angular momentum number, l, the shape of the orbital.
 Values allowed for l are 0-all positive integers less than or equal
to n-1.
 So orbitals for which n=2 can have one of two shapes corresponding to
l=0 and l=1. Depending on the value of l, an orbital is assigned a letter,
 0=s, 1=p, 2=d, 3=f.
More on quantum theory !
 Magnetic quantum number, (m)- is the
orientation of an orbital around the nucleus.
 Spin quantum number- has 2 possible values
(+1/2 or -1/2) which indicates the spin states of
an electron in an orbital.
Shapes of Orbitals
 http://www.youtube.com/watch?v=FfngYhl7ZmI
 S orbitals are spherical
 P orbitals are dumbbell shaped
 D orbitals are complex
 F orbits are more complex
 SO:
 n=1 one sublevel s
 n=2 two sublevels; s and p
 n=3 three sublevels; s,p and d
 n=4 four sublevels; s,p,d and f
 Seven f orbitals in each f sublevel
 Spin Quantum Number:
 Has only two possible values (+1/2,-1/2) which indicate the two
fundamental spin states of an electron in an orbital.
Rules Governing Electron
Configurations:
 Remember that real atoms are not built up by adding protons
and electrons one at a time.
 There are 3 rules:
 Aufbau principle: an electron occupies the lowest energy
orbital that can receive it.
 Pauli exclusion principle: no two electrons in the same atom
can have the same set of four quantum numbers.
 Hund’s rule: orbitals of equal energy are each occupied by
one electron before any orbital is occupied by a second
electron, and all electrons in singly occupied orbitals must
have the same spin state.
 These rules require placing as many unpaired electrons as
possible in separate orbitals in the same sublevel. So
electron-electron repulsion is minimized so that the
electron arrangements have the lowest energy possible.
 Applying Hund's rule shows that one electron will enter
each of the three p orbitals in a main energy level before a
second electron enters any of them.
 Hund’s rule: one electron will enter each of the three p
orbitals in a main energy level before a second electron
enters any of them.
Relative Energies of Orbitals
Orbital Notation
 An unoccupied orbital is represented by a line _____ with
the orbital’s name underneath the line.
 An orbital containing one electron is represented as ____.
 An orbital containing two electrons is represented as ____,
-¯
showing the electrons paired and with opposite spin states.
 So H and He are written like this:
H
He
He
-¯
1s
2nd period elements:
• In the first-period elements, hydrogen and helium, electrons
occupy the orbital of the first main energy level.
• According to the Aufbau principle, after the 1s orbital is
filled, the next electron occupies the s sublevel in the second
main energy level.
Elements of the 2nd period
Li - Ne
 The highest- occupied energy level is the electron-containing
main energy level with the highest principal quantum
number.
• The highest-occupied energy level is the electron-containing
main energy level with the highest principal quantum
number.
• Inner-shell electrons are electrons that are not in the highest-
occupied energy level.
Elements of the 3rd period:
Na - Ar
•
After the outer octet is filled in neon, the next electron
enters the s sublevel in the n = 3 main energy level.

Noble-Gas Notation
•
The Group 18 elements (helium, neon, argon, krypton,
xenon, and radon) are called the noble gases.
•
A noble-gas configuration refers to an outer main
energy level occupied, in most cases, by eight electrons.
Elements of the 4th period:
K - Kr
• The period begins by filling the 4s orbital, the empty orbital
of lowest energy.
• With the 4s sublevel filled, the 4p and 3d sublevels are the
next available vacant orbitals.
• The 3d sublevel is lower in energy than the 4p sublevel.
Therefore, the five 3d orbitals are next to be filled.
Elements of the 5th period:
Rb - Xe
• In the 18 elements of the fifth period, sublevels fill in a
similar manner as in elements of the fourth period.
• Successive electrons are added first to the 5s orbital, then to
the 4d orbitals, and finally to the 5p orbitals.
Elements of the 6th and 7th periods:
 6th period consists of 32 elements. To build up electron
configurations for elements of this period, electrons are
added first to the 6s orbital in Cesium, Cs, and barium, Ba.
Then, in lanthanum, La, an electron is added to the 5d
orbital.
 Cerium, Ce, the 4f orbitals begin to fill, giving cerium atoms
a configuration of [Xe]4f15d16s2. In the next 13 elements,
the 4f orbitals are filled. Next the 5d orbitals are filled and
the period is completed by filling the 6porbitals.
 Because the 4f and 5d orbital are very close in energy,
numerous deviations from the simple rules occur as these
orbital are filled.
 The 7th period is incomplete and consists largely of synthetic
elements, discussed in Nuclear Chemistry.