The building blocks of atoms

Every substance is made of tiny, fundamental building blocks called atoms. That's called the Atomic Hypothesis.

Nobel Prize-winning physicist Richard Feynman once said that if humankind had to give up all of its knowledge of science except for one fact, he'd keep the atomic hypothesis. He believed that from there - knowing that everything is made of atoms - we would be well on our way to regenerating all of our other knowledge about the universe.

Neutrons are particles that have the same mass (it's just slightly different) as protons, but have no charge. They contribute only in a minor way to the chemistry of atoms, but they do contribute to the mass, and are important for other reasons that we shall see in other sections.

Protons and neutrons compose the nucleus of an atom, which is "orbited" by electrons. While the nucleus accounts for practically all of the mass of an atom,

Elements

All of the stuff on (and of) Earth is composed of 90 naturally-occuring chemical elements. Two more occur naturally, but decompose so rapidly by radioactive decay that they aren't found. About 15 others can be made in laboratories, but they don't last long either — some on the order of just milliseconds.

What distinguishes one element from another is only the number of protons. The identity of an element doesn't depend in any way on the number of neutrons or electrons, though variances in those numbers do alter the properties of atoms.

We give each element a one- or two-letter abbreviation to make writing chemical equations more efficient. You should memorize the first ten elements now.

(Each element name in the table is a link to its Wikipedia page.)

The Roles of Electrons and Neutrons

Chemistry is mainly the science of how electrons behave in atoms and molecules, so we spend most of our time talking about them. However, an atom isn't an atom without protons and neutrons. Starting off, we'll think of atoms as charge-neutral, meaning that the sum of all charges, positive (the protons, charge = +1) and negative (the electrons, charge = -1) is zero. That means a neutral atom has to have an identical number of protons and electrons.

When the number of protons and electrons is not equal, we have an ion, a charged atom or molecule. For example, if a lithium (Li) atom (see above), which has three protons (that's what makes it a lithium atom), only has two electrons, its net charge is (+3) + (-2) = +1. If a fluorine (F) atom has 8 electrons and its complement of 7 protons (that's what makes it fluorine), its net charge is (+7) + (-8) = -1. We call negative and positive ions anions and cations, respectively.

There are specific conventions for labeling ions and isotopes. The table below is a summary. It's important that you learn how to label ions, especially. But be aware of isotope labeling, too.

Filling orbitals with electrons

In the section on electrons, we learned about the quantum nature of electrons bound to atoms by looking at hydrogen. We can generalize this theory to any atom by simply filling increasingly higher-energy orbitals with electrons, one at a time, as the atoms get bigger.

Learning to fill orbitals will provide the key to understanding the periodic table of the elements, and understanding the table is the key to much of what we'll be doing in chemistry as we go along.

It will help, as we go through the next bit of material, if you remember that (1) no two electrons in an atom can have the same set of quantum numbers (n, L, m_L and m_s) and (2) we "fill" orbitals from the lowest energy upward.

For the second atom, helium (He), with two protons, we place a second electron into the 1s orbital, noting that the two electrons must have different spins (there are only two kinds). This fills the 1s shell. It can't accept any more electrons.

For elements B, C, N, O, F, Ne, we fill in the three p sub-orbitals, which can each hold two spin-paired electrons for a total of six in the p-orbital. The electron configurations are:

B: 1s², 2s², 2p¹

C: 1s², 2s², 2p²

N: 1s², 2s², 2p³ ← 2p shell half filled

O: 1s², 2s², 2p⁴

F: 1s², 2s², 2p⁵

Ne: 1s², 2s², 2p⁶ ← 2p shell filled

You will be tempted to ask, with all those intersecting orbitals (which may not be drawn to scale, by the way), won't the electrons collide? Don't fall into that trap. Electrons don't act like particles when bound to atoms. There is avoidance of a sort, of course. Electrons must be spin paired; the spin pairs must live in different sub-orbitals; and negatively charged electrons repel each other in any situation.

Moving on to larger atoms, we start the process over again with n=3 (we're out of orbitals in the n=2 shell). Once we work through the 3s and 3p electrons (8 in all) we start filling the 3d orbitals ... except that there's a little kink. What really happens is that energy levels are always filled by the next highest level. It turns out that there's a bit of swapping of levels from what is expected. As an example, the electron configuration of Krypton is:

Kr: 1s², 2s², 2p⁶, 3s², 3p⁶, 4s², 3d¹⁰, 4p⁶

Notice that the 3d orbitals (5 of them) are filled between the 4s and 4p orbitals. We need to get to the bottom of that and find an easy way to remember it.

Sorting the energy levels

The diagonal rule

The diagonal rule makes remembering the energy level ordering easy. If you can remember how to write the levels in their logical order like this, then draw in the diagonal lines.

Now following the red arrows from upper right to lower left gives you the exact ordering. It's an easy way to remember.

Now given the number of electrons in an atom, you should be able to write the electron configuration.

Practice problems

Write the electron configurations of these atoms and ions. Remember that anions contain more electrons than a neutral atom and cations contain fewer. The configurations of the atoms/ions on the right use shorthand notation. Rollover or click the buttons to see the answer.