This chapter is intended as a review of concepts covered in more depth in general chemistry from the organic chemist's point of view. Bear in mind that many of the topics here are emphasized with an organic bent.
An atom consists of a nucleus of protons and neutrons surrounded by electrons. Each of the elements in the periodic table is classified according to its atomic number, which is the number of protons in that element's nucleus. Protons have a charge of +1, electrons have a charge of -1, and neutrons have no charge. Electrically, neutral atoms have the same number of electrons and protons, but they can have a varying number of neutrons. Within a given element, atoms with different numbers of neutrons are isotopes of that element. We will see that isotopes typically exhibit similar chemical behavior to each other.
Electrons have such little mass that they exhibit properties of both particles and waves. We know from Heisenberg's Uncertainty Principle that it is impossible to know the precise location of an electron. Despite this limitation, there are regions around the atom where the electron has a high probability of being found. Such regions are referred to as orbitals.
For isolated atoms (meaning non-bonded), electrons reside in the atomic orbitals of those atoms. Atomic orbitals are classified according to a set of four quantum numbers which describe the energy, shape, and orientation of the orbital.
Principle Quantum Number (n): Indicates how far the orbital is from the nucleus. Electrons are farther away for higher values of n . By Coulomb's Law we know that electrons which are closer to the positively charged nucleus are more powerfully attracted and thus have lower potential energies. Electrons of orbitals with higher values of n , being farther away from the nucleus, have greater potential energies. In a given atom, all the atomic orbitals with the same n are collectively known as a shell. n can take on integer values of 1 or higher (ex. 1, 2, 3, etc.).
Angular Momentum Quantum Number (l): Describes the shape of the orbital. The angular momentum number (or subshell) can be represented either by number (any integer from 0 up n -1) or by a letter ( s , p , d , f , g , and then up the alphabet), with 0 = s, 1 = p, 2 = d, and so on. For example:
when n = 1, l can only equal 0; meaning that shell n = 1 has only an s orbital (l = 0).
when n = 3, l can equal 0, 1, or 2; meaning that shell n = 3 has s , p , and d orbitals.
s orbitals are spherical, whereas p orbitals are dumbbell-shaped. d orbitals and beyond are much harder to visually represent.
Magnetic Quantum Number (m): Gives the orientation of the orbital in space; in other words, the value of m describes whether an orbital lies along the x -, y -, or z -axis on a three-dimensional graph, with the nucleus of the atom at the origin. m can take on any value from -l to l. For our purposes, it is only important that this quantum number tells us that for each value of n there may be up to one s -orbital, three p -orbitals, five d - orbitals, and so on: The s orbital (l = 0) has one orbital, since m can only equal 0. That orbital is spherically symmetrical about the nucleus.
Spin Quantum Number (s): Tells whether a given electron is spin up (+1/2) or spin down (-1/2). Because the Pauli Exclusion Principle tells us that no two electrons of an atom can have the same set of quantum numbers, each orbital is limited to holding two electrons at most.
It is often convenient to depict orbitals in an orbital energy diagram, as seen below in . Such diagrams show the orbitals and their electron occupancies, as well as any orbital interactions that exist. In this case we have the orbitals of the hydrogen atom with electrons omitted. The first electron shell ( n = 1) contains just the 1s orbital. The second shell ( n = 2) holds a 2s orbital and three 2p orbitals. The third shell ( n = 3) holds one 3s orbital, three 3p orbitals, and five 3d orbitals, and so forth. Note that the relative spacing between orbitals becomes smaller for larger n . In fact, as n gets large the spacing becomes infinitesimally small.
You will see such energy diagrams quite often in your continuing study of organic chemistry. Notice that all orbitals with the same n have the same energy. Orbitals with identical energies are said to be degenerate (not in the moral sense!). Electrons in higher level orbitals have more potential energy and are more reactive, i.e. more likely to undergo chemical reactions.
When an atom only contains a single electron, its orbital energies depend only on the principle quantum numbers: a 2s orbital would be degenerate with a 2p orbital. However, this degeneracy is broken when an atom has more than one electron. This is due to the fact that the attractive nuclear force any electron feels is shielded by the other electrons. s -orbitals tend to be closer to the nucleus than p -orbitals and don't get as much shielding, and hence become lower in energy. This process of breaking degeneracies within a shell is known as splitting. In general s orbitals become lowest in energy, followed by p orbitals, d orbitals, and so forth.