- Online Text
- 1. Introduction
- 2. Valence Electron Patterns and Lewis Structures
- 3. Ionic Bonds
- 4. Covalent Bonds and the Octet Rule
- 5. Polarity and Basic Lewis Structures
- 6. Advanced Lewis Structures
- 7. VSEPR Theory
- 8. Hybrid Orbitals
- 9. Intermolecular Forces
- 10. Physical Properties of Molecules
- 11. Conclusion
- 12. Further Reading
- Unit Guide (PDF)
Section 2: Valence Electron Patterns and Lewis Structures
Valence electrons, first introduced in Unit 4, Section 5, are the electrons in the outermost shells of the atoms, and they are the most reactive and the most important for understanding how atoms come together to form compounds. Most of the elements that we will consider in forming compounds are the main group elements, and for these elements their valence electrons are always in their outermost s and p orbitals. In Figure 5-2, all of the valence electrons for the common main group elements are represented as dots around the symbol. For hydrogen and helium, they can only have two electrons in their valence shell in that first row, but for the elements after that, we can see that the number of valence electrons ranges from 1 to 8. There is also a great periodic pattern that the elements in the same column or family on the periodic table always have the same number of valence electrons.
Figure 5-2. Valence Electrons for the First 26 Main Group Elements
Only the electrons in the outer, valence shell are represented. These range from one electron for the alkali metals to eight electrons for the noble gases. All of these electrons were found in the outermost s and p orbitals for each of these elements. Note that in the bottom row, the atomic number jumps from 20 to 31. The jagged line indicates where the transition metals would be located.
© Science Media Group.
Lewis Dot Structures
This notion of representing valence electrons as dots is a shorthand created by American chemist Gilbert Newton Lewis (1975–1946). Accordingly, these are called "Lewis dot structures" for atoms. Lewis was investigating how atoms bond, and in particular, the role of the electrons in the outermost shell. He found he could explain many compounds of two or more atoms by assuming that they could share electrons to give each atom a total of eight electrons in their outermost shell. (Figure 5-3)
Lewis didn't simply figure out how to diagram molecules with dots, he was also the one who explained the covalent bond, coined the term "photon," figured out how electrons pair together, and, among other things, he inspired Linus Pauling (1901–1994), another extraordinary chemist, to study the nature of chemical bonding. He was nominated 35 times for the Nobel Prize but never won; however, two of his students, Harold Urey (1893–1981) and Glenn Seaborg (1912–1999), won Nobel Prizes for their work on the discoveries of new elements and new isotopes.
Figure 5-3. Gilbert Newton Lewis
The chemist who discovered the covalent bond and created the Lewis dot structure method of depicting molecules.
© Bancroft Library, University of California.
As shown in Figure 5-2, Lewis dot structures are drawn by using the chemical symbol from the periodic table to represent the nucleus of an atom, and placing dots around the outside to represent the electrons in the valence shell. Note how once there are more than four electrons in the atom's Lewis structure, the electrons get paired together on the side of the atom. This is one of the key insights of Lewis: He realized that because orbitals hold two electrons, electrons on atoms like to pair together, even when they make bonds. With only four orbitals in the valence shell of most main group elements, electrons have to be paired once there are enough to go around.
However, Lewis structures for individual atoms aren't particularly useful, but this pairing of electrons will be very important when we get to larger multi-atomic structures. When electrons pair together on one atom, they are called "lone pair electrons" or a "lone pair." Note how neon, the unreactive noble gas has eight electrons arranged into four lone pairs. This is a simple representation of how its eight electrons are held in its four valence s and p orbitals. This number, "eight", is special in chemistry because so many elements have eight spaces in valence orbitals. We saw in Unit 4, Section 8 that atoms form ions in ways to try to gain or lose electrons to achieve these eight electrons that fill that outermost shell and therefore create a stable atom or ion. In the next few sections, we will learn how the Octet Rule will use this magical number, eight, to help us build most of the molecules in nature.