Supplemental material on chemical bonding

Electron orbitals - basically, the levels of energy for electrons in an atom - are important for understanding chemical bonds. Below is a figure that introduces you to a shorthand way of looking at orbitals that is helpful for understanding bonds. Each number (1-red circle, 2-green circles, and 3-orange circles) can be thought of as representing a level of energy. Each level can have 1 or more domains for pairs of electrons. The "s" orbitals are sometimes conceptualized as spheres and the "p" orbitals as dumbbells. The three p orbitals are conceptualized as lying along three axes at right angles to each other.

Each orbital can hold 2 electrons. The numbers below the circles indicate the order in which the orbitals fill as the atomic numbers get higher. Note that the s orbitals fill right away, but the p orbitals must fill in sequence starting with one electron each and then a second electron in each.

The element represented in this figure by the arrows in the 1s orbital is helium (He), which has two protons (the arrows are in opposite direction and represent the spin on each electron-spin isn't important for this class, but the opposite directions are handy for distinguishing the electrons). Because the 1s orbital is filled, the element has a neutral charge and is very stable. Electrons in the 2s orbital are not very tightly bound, so even though Be (beryllium, atomic number 4) would seem to be an equally stable element, it isn't-it loses the 2s electrons quite easily and becomes Be⁺². That's because the 2 orbital isn't completely filled. That orbital becomes filled at Ne (neon, atomic number 10). He and Ne are both so-called "noble gasses", which refers to the fact that they are completely unreactive, that is, their atoms do not bond with each other or with other atoms. What are the other noble gasses?

atoms that have filled orbitals are stable and inert

those with partially filled orbitals are reactive

Atoms "like" to be neutral, that is, have an equal number of electrons and protons, but they like even more to have full orbitals. The top figure below shows the orbitals for oxygen. Note the "lone" electrons in the 2p orbital. The 2p orbital is only partially filled, and a more stable state is for that orbital to be completely filled. So elemental oxygen would tend to acquire 2 more electrons in order to fill those orbitals. Although O^-2 doesn't actually occur as a free ion, it is often thought of that way because that describes how it bonds-it can easily form bonds with metals, for example, that lose two electrons.

The bottom picture below shows two oxygen atoms (orbitals on the blue oxygen have been reversed). Note that the electons in the half-empty 2p orbitals of the black oxygen compliment and fill the half-empty 2p orbitals of the blue oxygen. This is a very stable situation and, indeed, free oxygen tends to occur as O₂, the oxygen we breathe. This is a covalent bond, and is the stablest kind of bond.

Think about this: Why is O3-ozone-is such a dangerous pollutant in cities?

A covalent bond is when two atoms fill each others' orbitals.

Atoms can either gain or lose electrons to form ions. Here are sodium (Na) and chlorine (Cl), both of which readily ionize. Note that Na has a lone electron in an s orbital (3s). Such electrons are readily lost. In contrast, Cl has a lone electron in part of the 3p orbital; such an element will tend to pick up an extra electron. Losing electrons isn't possible because it can't lose just the lone electron. Orbitals lose electrons the same way they gain them, and the only way that Cl could achieve full orbitals by losing electrons would be to lose all 5 of the 3p electrons-that doesn't happen.

Ions like these will readily "share" electrons-the extra electron that Cl needs to fill the 3p orbital can come from the one that Na loses out of the 3s orbital. This is ionic bonding. Ionic bonds are not quite as stable as covalent bonds because the elements can occur as free ions. In contrast, O-2 never occurs as a free ion at surface conditions. The reasons are complicated and are related to the energy states of the various orbitals and within the orbitals at various levels of electron occupation.

Covalent bonds are usually much stronger than ionic bonds, but some bonds are intermediate in strength.

Two other types of bonds: van der Waals and metallic

van der Waals: best example is water, not the bonds between the H and O atoms -those bonds are covalent (use the circle models above to answer why) -but between water molecules. Below is a sketch of a water molecule (H₂O). Note that the H ions are on one side of the O atom. This means that, although the molecule as a whole is neutral, there is in fact a very slight negative charge on the O end of the molecule and a very slight positive charge on the H side of the molecule. This is one reason why water is such an extraordinary solvent for some types of substances. The charges can, in a sense, pull apart loosely bound solids such as salt (NaCl-see discussion of ionic bonds) by attracting positively or negatively charged ions.

Below is a figure showing three water molecules. This is more or less how water molecules arrange themselves in a liquid state; ice would be more regular. Note that the O ends of the molecules tend to lie next to the H ends.

metallic: Some metals, when they are in a pure or almost pure state (i.e., no other elements mixed in) form a kind of a "soup" of electrons; the bonds are strong but plastic, which is why many metals are malleable (for example, gold)