SAT Subject Test Chemistry
REVIEW OF MAJOR TOPICS
These are the skills that are usually tested on the SAT Subject Test in Chemistry. You should be able to …
• Define ionic and covalent bonds, and explain how they form.
• Identify the differences in the continuum that exists between ionic and covalent bonding.
• Describe the implications of the type of bond on the structure of the compound.
• Explain the implications of intermolecular forces and van der Waals forces.
• Explain how VSEPR and hybridization solve the need to comply with known molecular shapes.
This chapter will review and strengthen these skills. Be sure to do the Practice Exercises at the end of the chapter.
Some elements show no tendency to combine with either like atoms or other kinds of elements. These elements are said to be monoatomic molecules; three examples are helium, neon, and argon. A molecule is defined as the smallest particle of an element or a compound that retains the characteristics of the original substance. Water is a triatomic molecule since two hydrogen atoms and one oxygen atom must combine to form the substance water with its characteristic properties.
When atoms do combine to form molecules, there is a shifting of valence electrons, that is, the electrons in the outer energy level of each atom. Usually, this results in completion of the outer energy level of each atom. This more stable form may be achieved by the gain or loss of electrons or the sharing of pairs of electrons. The resulting attraction of the atoms involved is called a chemical bond. When a chemical bond forms, energy is released; when this bond is broken, energy is absorbed.
This relationship of bonding to the valence electrons of atoms can be further explained by studying the electron structures of the atoms involved. As already mentioned, the noble gases are monoatomic molecules. The reason can be seen in the electron distributions of these noble gases as shown in the following table.
The distinguishing factor in these very stable configurations is the arrangement of two s electrons and six p electrons in the valence energy level in five of the six atoms. (Note that helium, He, has only a single s valence energy level, which is filled with two electrons, making He a very stable atom.) This arrangement is called a stable octet. All other elements, other than the noble gases, have one to seven electrons in their outer energy levels. These elements are reactive to varying degrees. When they do react to form chemical bonds, usually the electrons shift in such a way that stable octets form. In other words, in bond formation, atoms usually attain the stable electron structure of one of the noble gases. The type of bond formed is directly related to whether this structure is achieved by gaining, losing, or sharing electrons.
Notice the reccurrence of the octet (8) of electrons in noble gases.
TYPES OF BONDS
When the electronegativity values of two kinds of atoms differ by 1.7 or more (especially differences greater than 1.7), the more electronegative atom will borrow the electrons it needs to fill its energy level, and the other atom will lend electrons until it, too, has a complete energy level. Because of this exchange, the borrower becomes negatively charged and is called an anion; the lender becomes positively charged and is called a cation. They are now referred to as ions, and the bond or attraction between them is called an ionic bond. These ions do not retain the properties of the original atoms. An example can be seen in Figure 11.
A 1.7 or greater electronegativity difference between atoms will essentially form an ionic bond.
These ions do not form an individual molecule in the liquid or solid phase but are arranged into a crystal lattice or giant ion-molecule containing many such ions. Ionic solids of this type tend to have high melting points and will not conduct a current of electricity until they are in the molten state.
Figure 11. Two Representations of the Ionic Bonding of LiF
When the electronegativity difference between two or more atoms is 0 or very small (not greater than about 0.4), the atoms tend to share the valence electrons in their respective outer energy levels. This attraction is called a nonpolar covalent bond. Here is an example using electron-dot notation and orbital notation:
Covalent bonds involve a sharing of electrons between atoms. Their electronegativity difference is between 0 to 0.5.
These covalent bonded molecules do not have electrostatic charges like those of ionic bonded substances. In general, covalent compounds are gases, liquids having fairly low boiling points, or solids that melt at relatively low temperatures. Unlike ionic compounds, they do not conduct electric currents.
When the electronegativity difference is between 0.4 and 1.6, there will not be an equal sharing of electrons between the atoms involved. The shared electrons will be more strongly attracted to the atom of greater electronegativity. As the difference in the electronegativities of the two elements increases above 0.4, the polarity or degree of ionic character increases. At a difference of 1.7 or more, the bond has more than 50% ionic character. However, when the difference is between 0.4 and 1.6, the bond is called a polar covalent bond. An example:
Notice that the electron pair in the bond is shown closer to the more electronegative atom. When these nonsymmetrical polar bonds are placed around a central atom, the overall molecule is polar. In the examples above, the chlorine (in HCl) and oxygen (in H2O) are considered the central atoms. Both the bonds and the molecules could be described as polar. Polar molecules are also referred to as dipoles because the whole molecule itself has two distinct ends from a charge perspective. Because of this unequal sharing, the molecules shown are said to be polar molecules, or dipoles. However, polar covalent bonds exist in some nonpolar molecules. Examples are CO2, CH4, and CCl4. (See Figure 12 .)
Figure 12 . Polar Covalent Bonds in Nonpolar Molecules
Polar covalent bonds have unequal sharing of electrons. Their electronegativity difference is between 0.4 and 1.6.
In all the examples in Figure 12 the bonds are polar covalent bonds, but the important thing is that they are symmetrically arranged in the molecule. The result is a nonpolar molecule.
In the covalent bonds described so far, the shared electrons in the pair were contributed one each from the atoms bonded. In some cases, however, both electrons for the shared pair are supplied by only one of the atoms. Two examples are the bonds in NH4+ and H2SO4. (See Figure 13 .)
Figure 13 . Covalent Bonds (both electrons supplied by one atom)
The formation of a covalent bond can be described in graphic form and related to the potential energies of the atoms involved. Using the formation of the hydrogen molecule as an example, we can show how the potential energy changes as the two atoms approach and form a covalent bond. In the illustration that follows, frames (1), (2), and (3) show the effect on potential energy as the atoms move closer to each other. In frame (3), the atoms have reached the condition of lowest potential energy, but the inertia of the atoms pulls them even closer, as shown in frame (4). The repulsion between them then forces the two nucleii to a stable position, as shown in frame (5).
In most metals, one or more of the valence electrons become detached from the atom and migrate in a “sea” of free electrons among the positive metal ions. The attractive force strength varies with the nuclear positive charge of the metal atoms and the number of electrons in this electron sea. Both of these factors are reflected in the amount of heat required to vaporize the metal. The strong attraction between these differently charged particles forms a metallic bond. Because of this firm bonding, metals usually have high melting points, show great strength, and are good conductors of electricity.
Metallic bonds are like positive ions in a “sea” of electrons.