• What holds atoms together?
• Review electron shell model of atom, valence electrons
• Define chemical bond
• Six types of bonds: ionic, metallic, covalent, polar, hydrogen, van der Waals. Be able to define each, provide examples, characterize their properties

• So far we discussed the nature of the atom in some detail, and have a qualitative sense of how it looks. However, a lone, non-interacting atom is rare. Most atoms are found in combination with others.
• After the big bang event the universe began to rapidly expand, and quite soon, within a few minutes, a major component was neutrons. Neutrons are not stable by themselves, so many of them split into protons and electrons, which formed a significant component of the unverse after about 10-15 minutes. It took about 100,000 years for the temperature of the universe to cool enough for the electrons to attach themselves to the protons and actually form atoms. So, about 100,000 years after the big bang, atoms became a significant component of the universe.
• This tells us that the energies of keeping electrons around a nucleus are much smaller than the energies associated with the nucleus or the formation of electrons to begin with.
• We live in the world of electrons, all our senses, and life itself, is manifest by variations in electronic interactions. Therefore life and humanity can only exist at the lower energy conditions in which electrons are bound to nucleii, i.e. Earth-like conditions and not Sun-like conditions.
• Electrons are the glue that holds groups of atoms together.

• Let's review the nature of the atom from the point of view of the electrons.
• The atom is mainly low-density space with a very small but dense nucleus that defines the center of the atom.
• Electrons are located around the nucleus.
• These electrons can be classified in terms of shells that correspond to the rows in the periodic table. Each shell can fit a certain number of electrons, depending upon how far away from the nucleus it is. A shell that is close to the nucleus can only contain a small number of electrons, otherwise the electrons are too close together, and electrostatic repulsive forces push them apart. periodic table
• Shells of electrons are most stable when they contain the maximum number of electrons that they can hold. On the one hand, if there are too few electrons, then the electrons are constantly whizzing about, trying to fill all of the available space. Therefore, they have high kinetic energy. On the other hand, if there are too many electrons then electrostatic repulsion takes over and pushes them apart.
• Define the electrons in an unfilled outer shell as valence electrons.
• Since these are the outermost electrons, these are the ones that are perturbed by bringing another atom close by. These are the ones that are involved in bonding.
• Define: a chemical bond is the result of a redistribution of electrons that leads to a more stable configuration between two or more atoms.

• We can classify chemical bonding into six major types: ionic, covalent, metallic, polar, hydrogen and van der Waals.

Ionic
• Ionic bonding occurs between a pair of atoms when one of the atoms gives up its valence electrons to the other. The result is that both atoms have filled shells. Both atoms also end up with a charge, one negative, and the other positive. We call the positive charged atom a cation, and the negatively charged one, an anion.

• A classic example of ionic bonding is between Na and Cl. Na is a silvery metal. It has 1 valence electron. Cl is a yellow-green gas, and it needs 1 electron to fill its valence shell. If you put the gas and the metal together, then they will burn as electrons are exchanged. The metal dissolves and the gas disappears. The ions now have opposite charges and are attracted to each other by electrostatic forces. They form a crystal with the rock salt structure.

• Ionic bonds commonly form between atoms of the 1st column and the 7th, and between the 2nd column and the 6th. E.g. NaCl, MgO
• Materials composed of ionic bonds have distinctive properties. They are strong when pushed together, but weak if cleaved or sheared.
• Once the bond breaks it is not easy to put back together. For instance, with NaCl. The Na is left 1 electron short, the Cl with one extra. They have no need to come close enough together to exchange electrons.
• Ionic crystals also dissolve easily in water, yet they have high melting points.
• So if you break a crystal apart, you can put it back together by dissolving it, or by heating it enough to excite the valence electrons off of the atoms.
• Ionic bonding is not only constrained to pairs of atoms. E.g. CaCl₂.

Metallic
• Metallic bonding occurs between atoms that have a small number of electrons in their valence shells. They give the electrons up, not just to one other atom, but to the complete group of atoms. We think of the electrons as becoming loose. Positively charged atoms sitting in a sea of electrons. We often call it electron-gas. E.g. sodium.

• This arrangement provides metals with many of their characteristic physical properties. The shiny luster is because with so many free electrons, and only two each in a given energy level, then the electrons form wide band of low energy, with the magnitude of the energy of the order of visible light. The electrons easily absorb light and re-emit it right back again, with no loss or gain of energy. This is the origin of the metallic luster.
• Metals can conduct electrons easily because the electrons are not localized in the crystal.
• Metals are malleable. This is because the atoms can easily rearrange themselves in the sea of electrons.
Covalent
• Covalent bonds often form between atoms with too many electrons in their valence shells to give away, but not enough to easily fill. Thus they share electrons with their neighbors, in such a way that including the shared electrons the shells are full.

• E.g. H₂. N₂, O₂, diamond.
• These are stronger bonds than either of the other two types. This is because the electrons are shared.
• The carbon bond usually forms covalent bonds. The possible arrangements are many. Life is based upon carbon-carbon bonding. The branch of science called organic chemistry is the study of carbon-based molecules.

• Your body is based upon carbon bonding. So the covalent bond is considered the most important bond with regards to life.
• Interestingly, Si, just above C in the periodic table, with its covalent bonding, is the basis for the computer industry.
• Another way to look at these things.
• Lets look at the electron densities to see another perspective on this bonding thing.

• Let’s make this clear. Nature does not form covalent bonds or ionic bonds. It just goes to the lowest energy arrangement that it can. We label it covalent or ionic.

Polar and hydrogen bonds

• This is a different kind of bond. It is not due to sharing of electrons, but results from the fact that some molecules distribute electrons in such a way that they have a region that seems charged. I.e. they become polar.
• E.g. water, H₂O. The side of the molecule with the hydrogens is slightly positive, the side with the O is slightly negative. The complete molecule is neutral. This means that the water molecule is stable, yet it can appear to be charged, depending upons its orientation. This is why water can dissolve so many things, e.g. water dissolves NaCl.

• Another example of a polar bond is called hydrogen bonding.
• Hydrogen bonding is very important. And is a part of most biological substances. Covalent molecules are often terminated by H atoms. These molecules are held together by  polar bonding of the H atoms.
• E.g. wood, plastics, silk, candle wax, DNA.
• Egg white is clear because of hydrogen bonding. But heat it up and break the bonds, and you end up with a white gelatinous solid.
• Polar and hydrogen bonds are weak.
Van der Waals
• Dynamics polar bonding. E.g. He₂

• E.g. clays. Sheets of strong covalent and ionic bonded atoms held together by weak Van der Waals forces. E.g. graphite. Like stacks of paper.