Bonding
Bonds are forces that hold atoms together. These bonded atoms form molecules. Bonds form between atoms to release energy. Atoms are more stable when they have less energy. Therefore, bonds with lower bond energies are generally more stable. Bonds also affect the chemical and physical properties of elements.
There are three types of bonds we will discuss here: ionic bonds, covalent bonds, and metallic bonds. (Note: hydrogen bonds are not actually molecular bonds, they are strong intermolecular forces between certain atoms in compounds.)
Ionic Bonds
Ionic bonds generally form between a metal and a nonmetal. These bonds have an electronegativity difference of 1.7 or greater (on Pauling's four-point scale; see Periodic Trends). In ionic bonding, ions are attracted to each other and form a crystal lattice. The force of attraction results from the transfer of electrons from the anion to the cation. In a crystal lattice, the smallest bonded molecule is called a formula unit. For instance, when Na+ and Cl- bond to form a crystal of sodium chloride, the formula unit is NaCl.
In the formation of ionic bonds, energy is released (so the process is exothermic). The more energy the process releases, the more stable is bond created. Also, the greater the charge difference between the atoms involved, the stronger the associated bond. Thus, a bond between barium and sulfur will be stronger than one between sodium and fluorine.
Compounds that are ionically bonded have high melting and boiling points. They are also brittle. A shift in the charged planes of the crystal causes shattering. Think about what would happen if the bottom row of the series of atoms below shifted to the right. Each atom would repel the similarly-charged atom across from itself, and this plane of the crystal would separate and shatter.
+ – + – + – + – + – + –
– + – + – + – + – + – +
Covalent Bonds
Covalent bonds generally form between two nonmetals. These bonds have an electronegativity difference of less than 1.7. The elements or molecules involved share electrons. Covalent bonds are weaker than ionic bonds.
Covalent bonds can be single, double, or triple, meaning that the particles involved share one, two, or three pairs of electrons. Most double and triple bonds are found with center atoms of carbon, oxygen, or nitrogen. Stronger bonds have a shorter bond length—the atoms are closer together. Double and triple bonds are stronger than singles, so they have shorter bond lengths. Triple covalent bonds are stronger than doubles which are in turn stronger than singles. Accordingly, bond energy (the energy required to break a bond) follows the same pattern.
One way to notate covalent bonds is to use a colon. The two dots represent the two electrons that the atoms are sharing. Thus, a single bond between two hydrogen atoms is written thus: H:H. A single bond can also be written with a dash: H–H. More dashes indicate more pairs of electrons involved. A double bond is shown thus: H=H.
Nonpolar covalent bonds
In nonpolar bonds, the atoms share electrons evenly. (The sharing creates no charged poles.) Nonpolar bonds have an electronegativity difference of less than .3.
Polar covalent bonds
In polar bonds, there is an uneven sharing of electrons. One atom gains a slight negative charge and the other becomes slightly positive. This effect is significant in the way polar molecules act with one another. Polar bonds have an electronegativity difference between .3 and 1.7 and are often gases at room temperature.
Metallic Bonds
Metallic bonds form between metal atoms. These bonds are the most different from the other types. When metallic elements bond to each other, they create a sea of electrons. The valence (outer shell) electrons of each atom are delocalized and move around the whole structure. For this reason, metals conduct electricity very well. Metals are also malleable and ductile. The strength of a bond increases with the number of valence electrons, so metals with more valence electrons, like lead and tin, form stronger bonds.
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