Chemistry Notes: Chapter 4
Light has wavelike characteristics and particle-like characteristics./
Wavelike properties:
1. Visible light travels through space as a wave of electromagnetic radiation just like x-ray, microwaves
and radio waves.
2. Light waves have a distinctive wavelength. It is the distance between corresponding points on
adjacent waves.
3. Frequency is the number of waves that pass a given point in a specific time. It is measured in Hertz.
One Hertz is one wave/second.
4. Frequency times wavelength gives us the speed of the wave. For light:
c = λ v where c is the speed of light.
Particle-like properties: Discovered because of the photoelectric effect which refers to the emission of
electrons from a metal when light shines on the metal. P. 93
1. Max Plank discovered that objects can emit energy in small units called quanta.
a. Quantum is the minimum quantity of energy that can be lost or gained by an atom.
b. E = hv where E is energy, v is the frequency of the radiation emitted and h is a constant
known as Planks constant. h = 6.625 x 10-34 Js.
c. In 1905, Einstein expanded on Planks theory by saying light behaves as a wave and a particle.
Photons—particle of electromagnetic radiation having zero mass and carrying a quantum
of energy with the energy depending on the frequency. Ephoton=hv
Ground State; lowest energy state of an atom.
Excited State—any state of energy higher than the ground state. As an excited atom returns to ground state, it
gives off energy in the form of light. (neon signs use this principle)
In 1813, Niels Bohr came up with a model of a hydrogen atom linking electrons with photon emission. He
stated that electrons circle the nucleus in well defined paths or “orbits”. Each orbit has a definite fixed amount
of energy with the lowest energy being nearest the nucleus. As excited electrons fall back to a lower energy
state, photons are emitted containing energy equal to the energy difference between the two orbits. However,
Bohr’s model did not explain atoms with more than one electron.
4-2 Quantum Model of the atom
In 1924 DeBroglie suggested that electrons be considered waves confined to a space around a nucleus. Thus,
they could be diffracted (bent) and experience interference when they overlap.
Heisenberg Uncertainty Principle (1927) it is impossible to determine simultaneously both the position and
velocity of an electron or any other particle because in the process of locating it, one knocks the electron off it’s
path. This principle still holds true today.
In 1926, Schrodinger developed an equation that treated electrons in atoms as waves.
Quantum Theory—Derived from Heisenberg’s uncertainty principle and Schrodinger’s wave equation. It
describes mathematically the wave properties of electrons and other very small particles.
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Quantum Numbers specify the properties of atomic orbitals and the properties of electrons in the orbitals ½.
There are four quantum numbers assigned to each electron in an atom.
1. Principle quantum number (n) indicates the main energy level occupied by the electron. 1,2,3,….
As “n” increases, so too does the distance from the nucleus.
2. Angular Momentum quantum number (l) indicates the shape of the orbital. The number of possible
shapes ie equal to “n”. If “n” is 3, there are three possible orbital shapes. Shapes are designated with letters: s,
p, d, and f
a. s orbital is spherical shaped.
b. p orbital is dumbbell shaped.
c. d and f shapes are very complex.
3. Magnetic Quantum Number (m) indicates the orientation of an orbital around the nucleus. Number
of total orbitals is equal to (n2)
a. Because “s” is spherical, it can have only one orientation.
b. “p” can have 3 extending along the x, y, and z axis.
c. “d” can have 5 orientations
d. “f” can have 7 orientations.
4. Spin Quantum ½ Number
(+½ or - ½) indicates the two fundamental spins of an electron in an orbital.
One orbital can hold two electrons and they must have opposite spins. Number of total electrons
is (and energy level is (2n2)
Section 4-3
Electron Configuration—arrangement of electrons in an atom—usually refers to the ground state of the
electron.
Three rules govern electron configurations.
1. Aufbau principle—electrons will always occupy the lowest energy orbital that can receive it.
2. Pauli exclusion principle—no two electrons in the same atom can have the same set of four quantum
numbers.
3. Hund’s rule—Orbits of equal energy are each occupied by one electron before they begin to double
up.
Orbital Notation—unoccupied orbital is represented by a line with the orbital name written underneath it.
Hydrogen

1s
Helium
Lithium
1s
1s

2s

Electron configuration notation—Number of electrons in a sublevel is shown by adding a superscript to the
sublevel designation. Oxygen—1s2 2s2 2p4
Elements of the second period—p. 110 ( 8 elements)
Elements of the third period—p. 111 (8 elements)
Elements of the 4th period—p. 112 (18 elements) fill the 4s, 3d and then 4p
Elements of the 5th period—(18 elements) fill the 5s, 4d and then 5p
Elements of the 6th period—(32 elements) fill the 6s, 4f, 5d, then 6p
Elements of the 7th period—not complete
Noble gas notation: designates the noble gas notation then adds what is extra so the inner shell electrons do not
have to be listed each time.
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Potassium: [Ar] 4s1
Cobalt [Ar] 4s2 3d7
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