1. Chem 110/Math 90
Atomic Models
Plum Pudding Model (1900s) Lord Kelvin/JJ Thomson—no nucleus!
Rutherford’s Model (early 1910s) Rutherford—atom has a nucleus!
Rutherford was one of the first to have a planetary model of the atom: the electron travels around the nucleus, somewhat like the way planets travel about the sun
PROBLEMS with this model:
No explanation for why the electrons
do not spiral into the nucleus—atoms should not exist!
Electrons in orbits of any energy should
omit light of all colors as they lose energy—but this is
not observed experimentally. (see Hydrogen line spectra)
This is a continuous spectrum
(light of all colors)—”WHITE LIGHT”
The Development of Atomic Structure

2. Atomic Models Line Spectra: Each element has its own line spectrum.
atoms of an element absorb energy (as light or heat)
as atoms lose energy, they emit light
the colors of the light seen can be separated by a prism
these separate colors make up the line spectra
The hydrogen line spectrum:
The helium line spectrum:
The neon line spectrum:

3. Atomic Models
Bohr’s Model (mid 1910s) Bohr said electrons travel in circular orbits. These orbits have fixed radii.
Electrons can “jump” from one orbit to another. When gaining energy, an electron jumps to an orbit farther from the nucleus. When losing energy, an electron jumps to an orbit closer to the nucleus.
The electrons of lowest energy stay at the orbit closest to the nucleus (they cannot get closer).

4. Bohr’s Model
"We are tracing the description of natural phenomena back to combinations of pure numbers which far transcends the boldest dreams of the Pythagoreans." -- Bohr
This model involves the idea of quantization (developed by Planck):
Energy levels (orbits) increase in quantized amounts of energy (discrete—not continuous—amounts of energy)
This means the light emitted from excited electrons are not continuous—only certain colors are seen (only QUANTA of energy can be absorbed and released)

5. Wave-Particle Duality
Bohr’s model did not explain why electrons do not spiral into the nucleus. (It makes the assumption that this does not occur).
It was accepted by physicists that light seemed to behave as both waves and as particles (called photons). This idea is called “Wave-Particle Duality.”
Louis de Broglie asked “If light can be wave-like and particle-like, why can’t matter also be wave-like and particle-like?”
This led to the concept that
electrons can behave as waves and as particles.
Prince Louis de Broglie
first got a history degree,
then went into science.
Won the Nobel Prize
in physics (1929).
His work helped shape
quantum mechanics.

6. Electrons as waves
Electrons behave like standing (stationary) waves (not traveling waves, like light)
Stationary waves are like the waves (vibrations) from plucking a guitar string.
There are only certain areas where the waves (electrons) are allowed. The 3-D space where electrons might be is described by orbitals.
This is one
kind of orbital
called a p-orbital

7. Schrödinger’s Model
The following model is now accepted:
The description of electrons as waves is given by a wave function.
A wave function can tell us what parts of the atom are likely to be occupied by electrons—this is called a probability density (an orbital!).
An orbital shows us the 3D space where an electron might be (>90% chance the electron is in this space)
(According to Werner Heisenburg, we cannot know the position and momentum of electrons simultaneously—thus, we can only speak of their location in terms of probability).
This is a probability
density (orbital).
The dots
show where
an electron is
likely to be
>90% of the time

8. Orbitals Wave functions determine the sizes and shapes of orbitals
Four common orbital types are: s, p, d, f
(theoretically, there are an infinite number of orbital types: g, h, I, etc…)
Let’s look at the difference between Bohr’s atom and Schrödinger’s:

9. Electronic structure of atoms
Electrons occupy different orbitals.
How are these orbitals arranged?
To know where they are likely to be, we give them an address called an “electron configuration”.
The electron configuration tells us where the electron is in terms of:
the energy level (shell)
the sub-energy level (subshell)
the orbital

10. Organization of electronic structure
The apartment analogy:
The shells (energy levels)
are like apartment floors
The subshells are like
apartments
The orbitals are like
rooms
The number of rooms
in each apartment depends
on the type of apartment
(s, p, d, f)
Each room only holds
two people (electrons) max

11. Electron Configurations
An electron configuration is a statement of how many electrons an atom has in each of its subshells.
Ex. Write the electron configuration for oxygen.

12. Steps to writing electron configurations
1. # of electrons in the atom of interest?
2. Write the subshells using the Aufbau (“building up”) principle
(Get the Aufbau order correct)
3. Fill in each subshell with the correct superscript (# of electrons in each subshell)

13. Electron Configurations
# of electrons in the atom of interest?
Oxygen has 8 electrons (atomic # = # of protons or # or electrons)
2. Write the subshells in order
of increasing energy, and put a slot for each orbital in that subshell
(s has one, p has three, d has five, etc…)

14. Order of subshells (in increasing energy)

15. So for oxygen, it would be: 2p ___ ___ ___ 2s ___ 1s ___
Fill in each orbital (slot) with the correct # of electrons.
Use arrows pointing in opposite directions.
4. What you have just done is a “orbital box diagram”
The electronic configuration (common form) is a shortcut:
All subscripts should add up to the number of electrons available

16. Interesting Links (to learn more about quantum mechanics)
Chem 110/Math 90
Interesting Links (to learn more about quantum mechanics)
These will be posted on our course website:
The Development of Atomic Structure