Know how Bohr's and Schrödinger's atomic models differ. Know the four quantum numbers and what each represents. State the Pauli Exclusion Principle. State Hund's rule. Write the electron configuration of any atom on the periodic table. Know how valence electrons are involved in chemical bonding. Know how they three basic types of chemical bonds are alike and how they are different. Know the seven diatomic molecules. Use the Octet Rule to describe bonding between atoms. Write the electron-dot diagram for any atom on the periodic table. Write the orbital-filling diagram for any atom on the periodic table.

Quantum Numbers:
four numbers used to describe the electrons in an atom.

The Bohr model was a one-dimensional model that used one quantum number to describe the electrons in the atom. Only the size of the orbit was important, which was described by the n quantum number. Schrödinger described an atomic model with electrons in three dimensions. This model required three coordinates, or three quantum numbers, to describe where electrons could be found.

The three coordinates from Schrödinger's wave equations are the principal (n), angular (l), and magnetic (m) quantum numbers. These quantum numbers describe the size, shape, and orientation in space of the orbitals on an atom.

1. Principal (shell) quantum number - n

Describes the energy level within the atom.
• Energy levels are 1 to 7
• Maximum number of electrons in n is 2 n ²

Describes the sublevel in n
• Sublevels in the atoms of the known elements are s - p - d - f
• Each energy level has n sublevels.
• Sublevels of different energy levels may have overlapping energies.

The momentum quantum number also describes the shape of the orbital.

• Orbitals have shapes that are best described as spherical (l = 0), polar (l = 1), or cloverleaf (l = 2).
• Orbitals even take on more complex shapes as the value of the angular quantum number becomes larger.

Describes the orbital within a sublevel
• s has 1 orbital
• p has 3 orbitals
• d has 5 orbitals
• f has 7 orbitals

Orbitals contain 1 or 2 electrons, never more.

m also describes the direction, or orientation in space for the orbital.

This diagram shows the three possible orientations of a p orbital - p_x, p_y, p_z. Atomic orbitals are studied in great detail in a college "theoretical inorganic chemistry" class.
If you are curious, study these websites:

This fourth quantum number describes the spin of the electron.
• Electrons in the same orbital must have opposite spins.
• Possible spins are clockwise or counterclockwise.

Rules governing the combinations of quantum numbers: The three quantum numbers (n, l, and m) are integers. The principal quantum number (n) cannot be zero. n must be 1, 2, 3, etc. The angular quantum number (l) can be any integer between 0 and n - 1. For n = 3, l can be either 0, 1, or 2. The magnetic quantum number (m) can be any integer between -l and +l. For l = 2, m can be either -2, -1, 0, +1, or +2. The spin quantum number (s) is arbitrarily assigned the numbers +1/2 and -1/2.

Pauli Exclusion Principle: No two electrons in an atom have the same set of four quantum numbers.

Hund's Rule: Electrons will enter empty orbitals of equal energy, when they are available.

Quantum Chemistry: Describes the way atoms combine to form molecules and the way molecules interact with one another, using the rules of quantum physics. One of the key insights in quantum chemistry is that, because an electron is not a classical particle located at a definite point in space, even a single electron can "surround" the nucleus of an atom, filling a volume roughly as big as the whole atom. Instead of thinking of electron shells neatly nested inside one another, it is better to visualize electrons in interpenetrating orbitals, like a lot of ripples on a pond. Each individual electron cloud extends down to "touch" the nucleus, and all the electrons in an atom come under the direct influence of the nucleus, although some are influenced more strongly than others. Since there is not enough room for all the orbitals to fit next to the nucleus, some of them are concentrated further out from the nucleus than others. Some arrangements of electrons in orbitals are more stable than others, and atoms will interact to reach these stable arrangements (the basis for chemical bonds).

Test Your Concept Understanding: When was Niels Henrik David Bohr born? When was Erwin Rudolf Josef Alexander Schrödinger born? What is the maximum number of electrons possible in the 6th electron energy level? How many sublevels are found in the 3rd electron energy level? How many orbitals are in the d sublevel of any energy level? What is the maximum number of electrons in the outer electron energy level of any atom? The atomic theory tells us that the fifth electron energy level has five sublevels. Why do we only have letters for four sublevels - s, p, d, f ? What are the possible numbers for m? What are the possible numbers for the fourth quantum number?

Chemistry Week 04 - Day 2 - 3

Electron configuration:

The shape of today's periodic table shows energy levels, sublevels, and orbitals.

• Except for helium, you can read the electron configuration directly from the table.
• If you mentally move helium beside hydrogen, above beryllium, it will be in place to read its electron configuration.
• Remember the energy overlaps in the d and f sublevels.

• There are a "few" exceptions, chromium being the first, but this class will not be concerned with the exceptions.

The electron configuration for chlorine is   1s² 2s² 2p⁶ 3s² 3p⁵

• The large numbers represent the electron energy level.
• The letters represent the electron energy sublevel.
• The superscript numbers indicate the number of electrons in the sublevel. Think of each square on the table as an electron. Counting the number of squares will give you the number of electrons in each energy level and sublevel.

  There are two ways to check an electron configuration: The last notation in the electron configuration should represent the location of the element on the periodic table. Example: the 3p5 in the electron configuration for chlorine above indicates its location as the fifth square in the p sublevel on the third row of the periodic table. The total of the superscripts in an electron configuration will equal the atomic number of the element. Example: the total of the superscripts in the electron configuration for chlorine above is 17, its atomic number.

Work these electron configuration practice problems. Ask your science facilitator if you have any questions about these problems.

Lab #1 Lab #2

In-class Assignment 043:
This assignment must be turned in by the end of class tomorrow to receive credit.

Review the parts of an essay.
Use either Google or Altavista for this assignment.

Choose one (1) of these:

1. Perform an Internet search for Schroedinger's Cat.
   Use this essay planning guide to write an essay about the topic.

2. Perform an Internet search for ERP Paradox.
   Use this essay planning guide to write an essay about the topic.

Research Links:

• Quatnum Numbers - University of Alberta
• Elementary Quantum Chemistry - University of California, Berkeley
• First Ionization Potential of the Elements - Massachusetts Institute of Technology
• Theory of Atoms in Molecules - McMaster University, Hamilton, Ontario
• The Quark Theory - Stanford Linear Accelerator Center
• Elementary Particles - University of Oregon
• Generic Naming Convention for the Chemical Elements - SweetHaven Publishing