Chemistry Tutorial

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Chapter 1


The atom is the building block of matter. It is not indivisible, though, as the Greek origin of its name indicates. Over the past two centuries, experimentalists have identified numerous particles that comprise an atom. For the purposes of chemistry, we are interested in three of these particles: the proton, the neutron, and the electron.

An atom consists of a central nucleus of protons and neutrons, and electrons that orbit the nucleus.

The proton has an electric charge of +1e, or 1.602 x 10-19 Coulombs (a Coulomb is a unit of electric charge. For convenience, the charges of subatomic particles are usually described as multiples of the value e, the charge of one electron, called the elementary charge.)

The neutron has no charge (it is electrically neutral.)

The electron carries a negative charge: -1e. Electrons are bound to the nucleus by their electromagnetic attraction to the positively-charged protons. An atom in its neutral state always has the same number of protons as electrons. The positive and negative charges cancel, and thus the net charge of the atom is zero.

The true nature of the electrons' movement around the nucleus is complex. As an introduction, it's useful to present the Classical Model of atomic structure. This is a simplified scenario in which we regard the electrons as discrete objects orbiting at known, fixed distances from the nucleus.


The electrons orbit only in certain "allowed" regions around the nucleus. The farther an electron orbits from the nucleus, the higher the energy associated with it. Through energy transfers during collisions with other particles or with light, it is possible for electrons to move up to orbitals of higher energy. In this case, we say that the atom is in an excited state. Here, we will only describe ground state electron configurations.

Each energy level, or shell, contains one or more orbitals. Each orbital contains one or more subshells. A subshell has a place for up to two electrons. Electrons "fill up" orbitals from the lowest energy up -- that is, the second orbital does not contain an electron unless the first orbital is already full.

Now, before you read any further, you might want to have the PERIODIC TABLE OF THE ELEMENTS handy, to follow along. This can get confusing!

Elements are classified by the number of protons in their nucleus
-- a number usually denoted by the letter Z. This value is called the atomic number of the element. The first element, Hydrogen (H), has an atomic number of Z=1. Its single electron travels around the proton in the first energy level, n=1 (where n is the integer rank number of the level and goes from 1 to infinity.) The n=1 energy level consists of a single 1S orbital, which is shaped like a sphere, centered on the nucleus (the nomenclature "S" is an historic artifact and is not significant.) Element with atomic number 2, Helium (He), contains two protons and two neutrons in its nucleus. Its two electrons fill the 1S orbital.

Beginning with Lithium (Li), with atomic number 3, there is not room in the n=1 energy level for all the electrons, so the n=2 level begins to fill. Lithium has two electrons in orbital 1S, and one electron in an n=2 orbital: 2S. Berillium (Be), number 4 in the Periodic Table, fills 2S with its two electrons.

Boron, though, begins to make use of the n=2 shell's more complex structure. The n=2 shell is constructed of two orbitals: 2S and 2P, and the 2P is further broken down into three subshells: 2Px, 2Py, and 2Pz. Each of the four subshells can contain two electrons, so in total the n=2 shell has room for eight electrons. The 2S is slightly lower in energy than the 2P's and thus fills first; the 2P's are all equal to each other in energy. The 2S orbital is shaped like a sphere; the three 2P's are each shaped like a dumbell, intersecting at their centers at right angles to each other.

Sodium (Na) is the first element to use the third energy level. N=3 is constructed the same way as the second -- with 3S filling before 3P.

We encounter 4S with Potassium (K) and Calcium (Ca), but at Scandium (Sc) the pattern changes. After the addition of the 4S electrons and before the addition of the 4P electrons, the sequence reverts to the third energy level to insert electrons in a 3D orbital. A D orbital can hold 10 electrons. After the 3D orbital is full, the sequence continues to 4P. (You might argue the existence of a contradiction here, since I earlier stated that lower energy levels fill first. The filling of the 3D orbital does not contradict this rule -- its energy is in fact lower than the 4P orbital energy, despite the numbering!)

Beginning with Lanthanide (La), the F orbital begins to fill, with 4F being filled after 6S.

The sequence of addition of electrons is as follows:
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6,
where the first number indicates the energy level, letter indicates orbital type, and second number indicates the number of electrons within each subshell.

Again, if you're having trouble grasping this pattern, study the PERIODIC TABLE. As you go from hydrogen down the chart, the groups 1 and 2 represent the filling of an S orbital. Groups 3 through 8 fill their P orbital. The transition elements fill the D orbital, and the Lanthanide and Actinide series fill the F orbital.


1. Element with electron configuration 1s2 2s2 2p3 is Nitrogen (N.) It has two electrons in 1s, two in 2s, and three in 2p (arbitrarily two in 2px and 1 in 2py.)

2. The electron configuration for Copper (Co) is: 1s2 2s2 2p6 3s3 3p6 4s2 3d7.

A chemist would shorten this notation to just "3d7" - calling Copper by the subshell of highest energy that contains any electrons.


Orbital classification alone will not give you the most useful picture of the way an atom is structured. In order to construct your Carbon atom model, keep in mind one additional point. An atom is not the compact, space-efficient object that high school chemistry books may illustrate. It is actually made up mostly of empty space. Take Hydrogen, for example. The volume of the atom itself (where an atom's "edge" is defined at the location of the outermost electron orbital) is approximately 2.8 x 109, or 2,800,000,000 times larger than its nucleus (the single proton.) WHY? With this scale in mind, regarding an atom as the miniature solar system pictured in textbooks is obviously a misleading view!


To calculate orbital radii for electrons in a Carbon atom, try using these approximations:

r = n2h2/[4Pi2mZe2]

    for n=1, and 

r = n2h2/[4Pi2m(Z-2)e2] 

    for n=2. WHY?  

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