Geos 306, Fall 2002, Lecture 2
The Nature of the Atom

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Important Points

• Definition of the atom: "That which cannot be divided" (Democritus, 530 BC)
• For our purposes we assume that the structure of the atom is composed of 2 parts, nucleus and electrons. The nucleus is made of protons and neutrons.
• Neutrons are the heaviest of these particles. They are unstable (1/2 life = 12 minutes) except in the presence of protons. They decay to form a proton and an electron.
• The proton has slightly less mass than a neutron, but more importantly, it has a positive charge. Protons can exist in close proximity with other protons only if neutrons are also present. Neutrons are the glue that holds a nucleus together. The number of protons in an atom is called the atomic number. This number defines an element. It also allows us to order the elements into a list. An example of such a list is the periodic table.
• The number of neutrons in a given atom is not fixed. The number of protons + the number of neutrons is called the mass number. Two atoms with the same atomic number but having different numbers of neutrons are called isotopes. Eg

¹²C (6 protons + 6 neutrons),
¹³C (6 protons + 7 neutrons),
¹⁴C (6 protons + 8 neutrons). One of the isotopes is always the most stable. Isotopes with odd mass numbers are the least stable. That means that ¹³C is less stable than ¹²C or ¹⁴C.

Associated with each proton is an electron. It has 1/2000^th the mass of a proton, and a charge that is equal in magnitude but opposite in sign to the charge of a proton.

Early models of the atom include the raisin muffin model, Rutherford's model and the Bohr model. The Bohr model led to the concept of quantized energy levels in an atom and wave/particle duality. Currently, the best solution to obtain the location and energies of electrons is provided by the Schrodinger equation.

The electrons are the glue that hold atoms together. As such, the nature of electrons and their interactions are the most important concept in chemistry. They are the basis for the periodic table.

Let’s examine the electronic structure of the atom. It can be described in terms of quantum numbers, with each electron associated with a unique set of quantum numbers. These are:

n = principal quantum number, (shell number or row number)
l = azimuthal quantum number (shape, s = sharp, p= principal, d = diffuse, f = fundamental, etc.)
m = magnetic quantum number, defines the orientation in a magnetic field, eg p_x, p_y, p_zetc.
s = spin quantum number, related to the axial spinning of an electron in the particle concept. We use the up and down arrow symbols.

• n can have any positive integer value, The number 1 is the first row of the periodic table, 2 is the second row etc.
• l can have values 0,1,2,…,n-1. So l = 0 is the s orbital, etc.
• m goes from –l to l. So we can have 1s,2s,2p_x,2p_y,2p_z, etc.
• s can only be –½ or ½.
• shapes s-orbitals p-orbitals d-orbitals f-orbitals
• Remember that this model is for the H atom, we hypothesize that the other atoms have these orbital shapes.
• Some examples.
O 1s², 2s²2p⁴
Mg 1s², 2s²2p⁶, 3s²
Si 1s², 2s²2p⁶, 3s²3p²
Fe 1s², 2s²2p⁶, 3s²3p⁶4s²3d⁶ Note that the energy levels do not increase simply by filling all n=1, then n=2, then n=3, etc. The order in which the energy levels fill can be found by examining the order of elements in the periodic table!
• Be able to make spin diagrams for various atom.

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Suggested Reading:
Nesse: pages 39-46
Amaldi, G. The nature of matter, Chicago Press, 1966.