Electron Structure of the Atom
Chapter 7

7.1 Electromagnetic Radiation and Energy

Electromagnetic Radiation
• EM Radiation travels through space as an oscillating waveform .
• EM Radiation travels through a vacuum at a constant speed of 3.00×10 8 m/s

Properties of EM Radiation
• Wavelength ( λ, measured in nm )
• Frequency ( υ, measured in Hertz, Hz )

Relationship between λ and υ

Electromagnetic Spectrum

Mathematical Relationships
υ = Frequency of the light (1/s, or Hz)
λ = Wavelength of light (nm or m) c = CONSTANT , Speed of light (3.00×10 8 m/s)

Mathematical Relationships
E photon
=hυ E photon
=(hc)/
υ = Frequency of the light (1/s, or Hz)
λ = Wavelength of light (nm or m) c = CONSTANT , Speed of light (3.00×10 8 m/s)
h = Planck’s Constant (6.626×10 -34 J×s)
E photon
= Energy of a single photon (J)

Example
• Assume we want to determine the frequency of orange light and the energy of a single photon of this light.
• Orange light = 600 nm = 6.00×10 -7 m
• υλ = c, therefore υ = c/λ
• 𝑣 =
• E
• E
3.00×10
8
6.00×10 −7 𝑚 photon photon 𝑚/𝑠 = 5.00×10
=3.31×10 -19 J
14 Hz
=hυ=(6.626×10 -34 J×s)(5.00×10 14 Hz)

PROBLEM
• Calculate the frequency and photon energy for an X-ray of wavelength 1.00 nm.
• X-Ray= 1.00 nm = 1.00×10 -9 m
• υλ = c, therefore υ = c/λ
• 𝑣 =
• E
3.00×10
8 photon 𝑚/𝑠
1.00×10 −9
= 3.00×10 17 Hz 𝑚
=hυ=(6.626×10 -34 J×s)(3.00×10 17 Hz)
• E photon
=1.99×10 -16 J

PROBLEM
• What color is laser with a frequency of
6.0×10 14 Hz?
• 𝑣λ = 𝑐 therefore
λ = 𝑐 𝑣
• λ =
3.00×10
8 𝑚/𝑠
6.0×10 14 𝐻𝑧
= 5.00×10
• 500 nm = Green Light
-7 m = 500 nm

Continuous vs. Line Spectra

7.2 The Bohr Model of the
Hydrogen Atom

Bohr Model of the Atom
• Propsed by Niels Bohr
• Explains the Emission
Spectrum of Hydrogen
• Relies of quantitized energy levels.
• Does not work for atoms with more than one electron.

7.3 The Modern Model of the
Atom

Orbitals and Orbits
• Bohr’s model had electrons orbit in tight paths , but this only worked for Hydrogen.
• Schrödinger expanded the model by using
3 dimensional orbitals

Energy Levels and Orbital Shape
• Electrons are still in quantitized energy levels.
• Orbitals of roughly the same size are in the same overarching, or principal , energy level .
• There are four ground state orbital geometries: s, p, d and f .

Naming Orbitals
• Orbitals are named for their principal energy level and their orbital geometry.
• The n=1 principal energy level has only one geometry, s.
• The n=2 principal energy level has two geometries, s and p.
• n=3 is composed of s, p, and d
• n=4 is composed of s, p, d and f.

Orbital Geometries

Orbital Diagrams

Rules for Filling in Orbitals
• Ground State Atoms have the same number of electrons as protons.
• Aufbau Principle – Start with the lowest energy level.
• Pauli Exclusion Principle – Max of two electrons in each orbital with opposite spins
• Hund’s Rule – Electrons are distributed in orbitals of the same energy as to maximize the number of unpaired electrons.

Sodium p= 11 e= 11
Example

Carbon
PROBLEM

Titanium
PROBLEM

Electron Configurations
• Orbital diagrams are informative but take a lot of space.
• Electron Configurations are a shorthand for these diagrams.
• Though they convey the same information, they do not show sublevel organization.

Sodium p= 11 e= 11
Example
Na 1s 2 2s 2 2p 6 3s 1

Nitrogen
PROBLEM

Iron
PROBLEM

7.4 Periodicity of Electron
Configuration

Periodic Table

Another Way to Look at It

7.5 Valance Electrons in the Main
Group Elements

Main Group Elements

Valance Electrons
• Valance Electrons are those electrons in the last filled principal energy level.
• Core Electrons are those below the valance level.
• Valance Electrons for Main Group Elements are those in the highest s and p orbitals .
• Main Elements in the same group have the same number of valance electrons.

7.6 Electron Configurations for
Ions

Sodium ion p= 11 e= 10
Example
Na 1s 2 2s 2 2p 6

Ion Electron Configurations
• Ion charges are as they are due to the role of orbitals.
• Ions are stable at 1+, 2+, or such because that gets the electron configuration to a completed principal energy shell (for main group elements).
• Na (1+) is isoelectronic with Neon (a completed n=2)

7.7 Periodic Properties of Atoms

Valance Electrons and Chemistry
• Valance electrons are the ones participating in chemical reactions.
• Compounds are stabilized by reaching a filled principal energy level.
• We will return to this next chapter.

Ionization Energy
• Ionization Energy , the amount of energy required to remove en electron from an gaseous atom (kJ/mol)
• The lower the ionization energy the more reactive a compound is.