A chemical bond is an attracting force that links two atoms together to form a molecule. There are several kinds of chemical bonds. Firstly covalent bonds will be discussed, the strong bonds that result from the sharing of electrons. Then will be examined other kinds of interactions, including hydrogen bonds that are weaker than covalent bonds but enormously important to biology. Finally, will be considered ionic bonding which is a consequence of the loss or gain of electrons by atoms.
Covalent bonds consist of shared pairs of electrons
When two atoms attain stable electron numbers in their outermost shells by sharing one or more pairs of electrons, a covalent bond forms. Consider two hydrogen atoms in close proximity, each with a single unpaired electron in its outer shell. Each positively charged nucleus attracts the other atom’s unpaired electron but this attraction is balanced by each electron’s attraction to its own nucleus. Thus, two unpaired electrons become shared by both atoms filling the outer shells of both of them. Two atoms are, thus, linked by a covalent bond and a hydrogen gas molecule (H2) is formed. A molecule made up of more than one type of atom is called a compound. A molecular formula uses chemical symbols to identify the different atoms in a compound and subscript numbers to show how many of each type of atoms are present. Thus, the formula for sucrose—table sugar—is C12H22O11. Each compound has a molecular weight (molecular mass) that is the sum of the atomic weights of all atoms in the molecule. Looking at the Periodic table in Annex 3 you can calculate the molecular weight of table sugar to be 342. Molecular weights are usually related to a molecule’s size. A carbon atom has a total of six electrons; two electrons fill its inner shell and four are in its outer shell. Because its outer shell can hold up to eight electrons, carbon can share electrons with up to four other atoms—it can form four covalent bonds. When an atom of carbon reacts with four hydrogen atoms, a molecule called methane (CH4) forms. Thanks to electron sharing, the outer shell of methane’s carbon atom is filled with eight electrons and the outer shell of each hydrogen atom is also filled. Four covalent bonds—each consisting of a shared pair of electrons—hold methane together. Table shows the covalent bonding capacities of some biologically significant elements.
Orientation of covalent bonds
Covalent bonds are very strong. The thermal energy that biological molecules ordinarily have at body temperature is less than 1 percent of that needed to break covalent bonds. Thus biological molecules, most of which are put together with covalent bonds, are quite stable. This means that their three-dimensional structures and the spaces they occupy are also stable. A second property of covalent bonds is that, for a given pair of atoms, they are the same in length, angle and direction, regardless of the larger molecule of which the particular bond is a part. The four filled orbitals around the carbon nucleus of methane, for example, distribute themselves in space so that the bonded hydrogens are directed to the corners of a regular tetrahedron with carbon in the centre. The three-dimensional structure of carbon and hydrogen is the same in a large, complicated protein as it is in the simple methane molecule. This property of covalent bonds makes the prediction of biological structure possible. Although the orientations of orbitals and the shapes of molecules differ, depending on the types of atoms involved and how they are linked together, it is essential to remember that all molecules occupy space and have three-dimensional shapes. The shapes of molecules contribute to their biological functions.
Unequal sharing of electrons.
If two atoms of the same element are covalently bonded, there is an equal sharing of the pair(s) of electrons in the outer shell. However, when two atoms are of different elements, the sharing is not necessarily equal. One nucleus may exert a greater attractive force on the electron pair than the other nucleus, so that the pair tends to be closer to that atom. The attractive force that an atom exerts on electrons is its electronegativity. It depends on how many positive charges a nucleus has (nuclei with more protons are more positive and thus more attractive to electrons) and how far away the electrons are from the nucleus (closer means more electronegativity).
The closer two atoms are in electronegativity, the more equal their sharing of electrons will be. Looking at the table electronegativities of some elements important in biological systems, it is obvious that two oxygen atoms, both with electronegativity of 3.5, will share electrons equally, producing what is called a nonpolar covalent bond. So will do two hydrogen atoms (both 2.1). But when hydrogen bonds with oxygen to form water, the electrons involved are unequally shared: they tend to be nearer to the oxygen nucleus because it is the most electronegative of two. The result is called a polar covalent bond. Because of this unequal sharing of electrons, the oxygen end of the hydrogen–oxygen bond has a slightly negative charge (symbolized D– and spoken as “delta negative,” meaning a partial unit of charge) and the hydrogen end is slightly positive (D+). The bond is polar because these opposite charges are separated at the two ends or poles of the bond. The partial charges that result from polar covalent bonds produce polar molecules or polar regions of large molecules. Polar bonds greatly influence the interactions between molecules containing them.
Hydrogen bonds may form within or between atoms with polar covalent bonds
In liquid water, the negatively charged oxygen atom of one water molecule is attracted to the positively charged hydrogen (D+) atoms of another water molecule. (Remember, negative charges attract positive charges.) The bond resulting from this attraction is called a hydrogen bond. Hydrogen bonds are not restricted to water molecules. They may form between an electronegative atom and a hydrogen atom covalently bonded to a different electronegative atom. A hydrogen bond is a weak bond; it has about one-tenth (10%) the strength of a covalent bond between a hydrogen atom and an oxygen atom. However, where many hydrogen bonds form, they have considerable strength and greatly influence the structure and properties of substances. Hydrogen bonds also play important roles in determining and maintaining the three-dimensional shapes of giant molecules such as DNA and proteins.
Ionic bonds form by electrical attraction
When one interacting atom is much more electronegative than the other, a complete transfer of one or more electrons may take place. Consider sodium (electronegativity 0.9) and chlorine (3.1). A sodium atom has only one electron in its outermost shell; this condition is unstable. A chlorine atom has seven electrons in its outer shell—another unstable condition. Since the electronegativities of these elements are so different, any electrons involved in bonding will tend to be much nearer to the chlorine nucleus—so near, in fact, that there is a complete transfer of the electron from one element to the other. This reaction between sodium and chlorine makes the resulting atoms more stable. The result is two ions. Ions are electrically charged particles that form when atoms gain or lose one or more electrons.
The sodium ion (Na+) has a +1 unit charge because it has one less electron than it has protons. The outermost electron shell of the sodium ion is full, with eight electrons, so the ion is stable. Positively charged ions are called cations.
The chloride ion (Cl–) has a –1 unit charge because it has one more electron than it has protons. This additional electron gives Cl– a stable outermost shell with eight electrons. Negatively charged ions are called anions. Some elements form ions with multiple charges by losing or gaining more than one electron. Examples are Ca2+ (calcium ion, created from a calcium atom that has lost two electrons) and Mg2+ (magnesium ion). Two biologically important elements each yield more than one stable ion: Iron yields Fe2+ (ferrous ion) and Fe3+ (ferric ion) and copper yields Cu+ (cuprous ion) and Cu2+ (cupric ion). Groups of covalently bonded atoms that carry an electric charge are called complex ions; examples include NH4+ (ammonium ion), SO42– (sulfate ion), and PO43– (phosphate ion). The charge from an ion radiates from it in all directions. Once formed, ions are usually stable and no more electrons are lost or gained. Ions can form stable bonds resulting in stable solid compounds such as sodium chloride (NaCl) and potassium phosphate (K3PO4).
Ionic bonds are bonds formed by electrical attraction between ions bearing opposite charges. In sodium chloride—familiar to us as table salt—cations and anions are held together by ionic bonds. In solids, the ionic bonds are strong because the ions are close together.
However, when ions are dispersed in water, the distance between them can be large; the strength of their attraction is thus greatly reduced. Under the conditions that exist in the cell, an ionic attraction is less than one-tenth as strong as a covalent bond that shares electrons equally. Not surprisingly, ions with one or more units of charge can interact with polar molecules as well as with other ions.
Such interaction results when table salt or any other ionic solid, dissolves in water: “shells” of water molecules surround the individual ions, separating them. The hydrogen bond described earlier is a type of ionic bond because it is formed by electrical attraction. However, it is weaker than most ionic bonds because it is formed by partial charges (D+ and D–) rather than by whole-unit charges (+1 unit, –1 unit).
Polar and nonpolar substances interact best among themselves
“Like attracts like” is an old saying and nowhere is it more true than in polar and nonpolar molecules which tend to interact with their own kind. Just as water molecules interact with one another through polarity-induced hydrogen bonds, any molecule that is itself polar will interact with other polar molecules by weak (D+ to D–) attractions in hydrogen bonds. If a polar molecule interacts with water in this way, it is called hydrophilic (“water-loving”). What about nonpolar molecules? For example, carbon (electronegativity 2.5) forms nonpolar bonds with hydrogen (electronegativity 2.1). The resulting hydrocarbon molecule— that is, a molecule containing only hydrogen and carbon atoms—is nonpolar, and in water it tends to aggregate with other nonpolar molecules rather than with polar water. Such molecules are known as hydrophobic (“waterhating”) and the interactions between them are called hydrophobic interactions. It is important to realize that hydrophobic substances do not really “hate” water; they can form weak interactions with it (recall that the electronegativities of carbon and hydrogen are not exactly the same). But these interactions are far weaker than the hydrogen bonds between the water molecules and so the nonpolar substances keep to themselves.