SAT Chemistry Bonding - Types Of Bonds
SAT Chemistry Bonding - Types Of BondsIonic 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.
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.
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:
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 mol-ecule is polar. In the examples above, the chlorine (in HCI) and oxygen (in H2O) are consid-ered 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.)
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 H2S04. (See Figure 13.)
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.
INTERMOLECULAR FORCES OF ATTRACTION
The term intermolecular forces refers to attractions between molecules. Although it is proper to refer to all intermolecular forces as van der Waals forces, named after Johannes van der Waals (Netherlands), this concept should be expanded for clarity.
One type of van der Waals forces is dipole-dipole attraction. It was shown in the discussion of polar covalent bonding that the unsymmetrical distribution of electronic charges leads to positive and negative charges in the molecules, which are referred to as dipoles. In polar molecular substances, the dipoles line up so that the positive pole of one molecule attracts the negative pole of another. This is much like the lineup of small bar magnets. The force of attraction between polar molecules is called dipole-dipole attraction. These attractive forces are less than the full charges carried by ions in ionic crystals.
London Dispersion Forces
Another type of van der Waals forces is called London dispersion forces. Found in both polar and nonpolar molecules, it can be attributed to the fact that a molecule/atom that usually is nonpolar sometimes becomes polar because the constant motion of its electrons may cause uneven charge distribution at any one instant. When this occurs, the molecule/atom has a temporary dipole. This dipole can then cause a second, adjacent atom to be distorted and to have its nucleus attracted to the negative end of the first atom. London dispersion forces tire about one-tenth the force of most dipole interactions and are the weakest of all the electrical forces that act between atoms or molecules. These forces help to explain why nonpolar substances such as noble gases and the halogens condense into liquids and then freeze into solids when the temperature is lowered sufficiently. In general, they also explain why liquids composed of discrete molecules with no permanent dipole attraction have low boiling points relative to their molecular masses. It is also true that compounds in the solid state that are bound mainly by this type of attraction have rather soft crystals, are easily deformed, and vaporize easily. Because of the low intermolecular forces, the melting points are low and evaporation takes place so easily that it may occur at room temperature. Examples of such solids are iodine crystals and moth balls (paradichlorobenzene and naphthalene).
A proton or hydrogen nucleus has a high concentration of positive charge. When a hydrogen atom is bonded to a highly electronegative atom, its positive charge will have an attraction for neighboring electron pairs. This special kind of dipole-dipole attraction is called a hydrogen bond. The more strongly polar the molecule is, the more effective the hydrogen bonding is in binding the molecules into a larger unit. As a result the boiling points of such molecules are higher than those of similar nonpolar molecules. Good examples are water and hydrogen fluoride.
Studying Figure 14 shows that in the series of compounds consisting of H2O, H2S, H2Se, and H2Te an unusual rise in the boiling point of H2O occurs that is not in keeping with the typical slow increase of boiling point as molecular mass increases. Instead of the expected slope of the line between H2O and H2S, which is shown in Figure 14 as a dashed line, the actual boiling point of H2O is quite a bit higher—100°C. The explanation is that hydrogen bonding occurs in H2O but not to any significant degree in the other compounds.
This same phenomenon occurs with the hydrogen halides (HF, HCl, HBr, and HI). Note-in Figure 14 that hydrogen fluoride, HF, which has strong hydrogen bonding, shows an unexpectedly high boiling point.
Hydrogen bonding also explains why some substances have unexpectedly low vapor pressures, high heats of vaporization, and high melting points. In order for vaporization or melting to take place, molecules must be separated. Energy must be expended to break hydrogen bonds and thus break down the larger clusters of molecules into separate molecules. As with the boiling point, the melting point of H20 is abnormally high when compared with the melting points of the hydrogen compounds of the other elements having six valence electrons, which are chemically similar but which have no apparent hydrogen bonding. The hydrogen bonding effect in water is discussed on pages 181-183.
DOUBLE AND TRIPLE BONDS
To achieve the octet structure, which is an outer energy level resembling the noble gas configuration of eight electrons, it is necessary for some atoms to share two or even three pairs of electrons. Sharing two pairs of electrons produces a double bond. An example:
It can be assumed from these structures that there is a greater electron density between the nuclei involved and hence a greater attractive force between the nuclei and the shared electrons. Experimental data verify that greater energy is required to break double bonds than single bonds, and triple bonds than double bonds. Also, since these stronger bonds tend to pull atoms closer together, the atoms joined by double and triple bonds have smaller inter-atomic distances and greater bond strengths, respectively.