Pattern Recognition

Rules of thumb (ROTs) are useful devices to help students remember certain facts and to increase their ability to apply general principles. Generalizations may be made and expressed as ROTs. As generalizations, ROTs have exceptions but provide a core upon which students can build, through experience. These ROTs are designed to help students recognize situations within which they can work successfully.

1. Representative element metals and nonmetals generally tend to bond ionically, forming ionic solids.

2. Nonmetals bonding with nonmetals usually bond covalently, forming molecular solids, molecular liquids, and, in some cases, network solids.

3. When an electronegativity difference is present between atoms forming a covalent bond, the resulting bond is polar.

4. An electronegativity difference between bonded atoms producing more than 50% ionic character is the arbitrary demarcation between ionic and covalent bonds.

Common Student Misconceptions

1. "Bonding must be either ionic or covalent."
   It is a common student misconception that a bond between two atoms, A-B, is either purely covalent or purely ionic. No compound is 100% ionic. If the bond involves the same atoms (a homonuclear bond, A-A) then the bond must be 100% covalent because neither atom has the ability to attract the electron pair more strongly than the other. However, if the bond involves different atoms (a heteronuclear bond, A-B) the bond will have mixed covalent and ionic character. This means there will be a percent ionic character. Thus, except when the two atoms that are bonded are the same element (for example, two oxygen atoms), a bond is always partially covalent, partially ionic. The reason for this is that an electron is never completely transferred from one atom to another. The electron is shared rather than completely transferred. The sharing is a matter of degree-the concept of a polar bond. The best way to teach bonding is to show that there is a gradual progression from 100% pure covalent bond (homonuclear) to one that is about 98% ionic.

2. "Bond energies can be reliably used to calculate heats of reaction."
   This value may differ from ΔH determined from heats of formation since the bonding energy tables are derived from averages of bond energies. It is therefore best to avoid using bond energies to calculate ΔH.

3. "Mathematical or mental constructs such as electron clouds represent something solid."
   When solid models are used to illustrate atomic scale events, it is possible to produce this misconception among students. It is quite natural for students to develop this picture since both solid models and pictures resemble solids. Caution students about taking an overly-literal view of such models.

4. "Intermolecular bonds are the same as intramolecular bonds."
   There's quite a difference between these two types of bonds, despite their "sound-alike" nature. For example, liquid nitrogen has a low boiling point (-147 ∞C) due to relatively weak intermolecular van der Waals forces. Yet temperatures of several thousand degrees Celsius do not cause appreciable decomposition of N2 molecules because the nitrogen atoms in the N2 molecule are bonded with a triple covalent intramolecular bond. This is a strong, primary chemical bond whereas van der Waals forces are quite weak intermolecular forces that require relatively little energy to overcome.

5. "Covalent bonds must be weak because covalent compounds are generally soft with low melting points (< 300 ∞C)."
   Actually, this is a case of confusing intermolecular with intramolecular bonding. See the discussion above for Item 4.

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