Chemistry
Chemistry 8: 1-3
WebLecture
Valence Electrons and Lewis Diagrams
Outline
- Representing Bonds in Molecules
- Practice with Concepts
- Discussion Questions
- Optional Website Reading
Representing Bonds in Molecules
An electron configuration is completely balanced and at its lowest possible level when it mimics the configuration of a noble gas (Group 8A). A non-noble gas element can achieve this kind of configuration by becoming an ion (losing or gaining electrons), or by forming a bond with other atoms and sharing their electrons. We've discussed at length how ions form, now lets look at how bonds form to create complete configurations.
Valence Electrons
An electron configuration is completely balanced and at its lowest possible level when it mimics the configuration of a noble gas (Group 8A). A non-noble gas element can achieve this kind of configuration by becoming an ion (losing or gaining electrons), or by forming a bond with other atoms and sharing their electrons. We've discussed at length how ions form, now lets look at how bonds form to create complete configurations.
Not all the electrons in a particular atom can be shared. The electrons in lower levels are shielded by the outer electrons and closely held by the positive charge on the nucleus. So only the electrons in the outer levels can be valence or bond-capable electrons.
Rules for figuring out whether electrons are valence electrons.
- Any s electrons in the highest energy level are valence electrons.
- In elements with no p electrons (groups 1A, 2A, and 1B to 8B), the d electrons of the energy level immediately below the s level electrons are unshielded from outside influences, so some of these can actually take part in bonding. In this case, the outermost energy level s electrons and the most immediate lower energy level d electrons are the valence electrons.
- In elements with p electrons (groups 3A to 7A), the d electrons of the immediately lower level are shielded so that they cannot take part in chemical bond formation; only the s and p electrons of the outermost energy level are valence electrons. The d electrons are considered part of the core.
- Electrons in the lower energy levels that are part of the "last" noble gas configuration are core electrons, and do not take part in bonding.
Lewis dot symbols
| We can represent the electron configuration of the valence electrons by dots placed around the chemical symbol. For neon, with a full valence shell, there are pairs of electrons on all four sides. |
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| When there are an even number of electrons, the configuration reflects Hund's rule that unpaired electrons will occupy the available shells. In oxygen, there is one full 2s shell (one pair of electrons), on full two p shells (a second pair of electrons), and each of the two remaining electrons occupy one of the remaining p shells (single electrons shown). |
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Octets
| Chemical bonds attempt to create complete valence shells. Most molecules arrange the atoms so that there are 8 electrons around each atom (2 around hydrogen and helium). When we represent the water molecule with a Lewis dot diagram, each hydrogen has two dots, and the oxygen has 8 (a complete octet): |
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| We can also represent bonds by dashes, so the water molecule could also be shown as |
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| In some cases, the octet only forms if atoms share two electrons each, as with carbon dioxide. The double bonds can be represented either by two pairs of dots, or by lines. |
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Bond Types
All sharing is really a tug-of-war situation: the electron is drawn to two (or more) positively-charged nuclei. The electron will spend most of its time closer to the nucleus with the stronger electromagnetic field. So we have two extremes for electron sharing: either the nuclei attract the electron equally and it is shared perfectly between them (a covalent bond), or the electron is attracted totally to one atom and not shared at all by the other, creating a charge difference that holds the atoms together (ionic bond). Note that in the later case, the electron is not shared: what holds the two atoms together is the net difference in charge.
Pure covalent bonds occur only between atoms of the same element. O2 has a double covalent bond, O=O.
Ionic bonds occur when an element with low ionization energy (which means that it loses electrons easily) combines with an element that has high electron affinity (which means that it grabs electrons whenever it can). Elements of the 1A and 2A group have low ionization energies; elements of the 6A and 7A groups have high electron affinity and high ionization energies so that bonds between metals and halogens, such as NaCl, tend to be ionic.
It is important to realize atoms that ionic and covalent bonds are not opposites but the extreme ends of a range of possible sharing options between two atoms.
Ionic Bonds
Two steps are required for the formation of an ionic bond from neutral atoms. First, one component must lose one or more electrons, becoming a cation with a net positive charge (more protons than electrons). The other component must attract and hold the electrons, becoming an anion with a net negative charge (more electrons than protons). Each of these operations requires the input of energy, so they are both endothermic reactions.
The opposite charges draw the cation and anion together. This is an energy-releasing action, or exothermic reaction. If the energy released by the combining of the ions is greater than the energy required to form the cation and anion, the resulting ionic bond will hold the two ions together until some other electrical attraction or collision forces them apart.
To determine the total energy required to create the ions, we add the following together:
- the ionization energy of the element that becomes the cation
- the electron affinity of the element that becomes the anion
and subtract the ion-pair formation energy. If the amount is positive (endothermic), the ionic compound will not form.
Now, most ionic compounds do not form distinct molecules. Rather, they form extended solids, with each cation attracting several anions. NaCl, for example, forms a lattice in three dimensions, much like a 3-dimensional checkerboard, with sodium ions where the black squares would go and chlorine ions where the red squares would go. There is no reason that the edge chlorine should not attract other sodium ions; the "edge" of the crystal is arbitrary. So each chlorine is attracted by six sodium ions (remember to think in three dimensions! There are layers of sodium/chlorine lattices above and below this one).
| Cl | Na | Cl | Na |
| Na | Cl | Na | Cl |
| Cl | Na | Cl | Na |
| Na | Cl | Na | Cl |
Because each anion is held in place by multiple cations, the lattice bond is stronger than the simply ionic bond of the ion-pair. The actual lattice bond energy depends not only on the elements involved, but also on how they are packed. We will look at different crystal lattice possibilities in a later chapter.
Practice with the Concepts
Discussion Questions
- Which is stronger, an ionic bond or a covalent bond?
- Why does knowing bond type help you predict bond strength?
- How are the electric field attractions in ionic bonds different from the electric field attractions in covalent bonds?
- What happens to valence electrons that do not participate in bonds?
Optional Readings
- Learn more about drawing Lewis diagrams at the University of Calgary's Chemistry department Chemical Bonding web page.
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