Chemical Bonds

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Chapter: Anatomy and Physiology for Health Professionals: Levels of Organization : Chemical Basics of Life

Atoms can bond with other atoms by using chemical bonds that result from interactions between their electrons.

Chemical Bonds

Atoms can bond with other atoms by using chemical bonds that result from interactions between their electrons. During this process, the atoms may gain, lose, or share electrons. Chemically inactive atoms are known as inert atoms. An example of a chemical that is made up of inert atoms is helium. Atoms that either gain or lose electrons are called ions. These atoms are electrically charged. An example of an electrically charged atom or ion is sodium. The three important types of chemical bonds are ionic, covalent, and hydrogen.

Ionic Bonds

Ionic bonds form between ions. They are chemical bonds between atoms that form because of the transfer of electrons. The atom gaining one or more electrons is referred to as the electron acceptor. Ions that acquire a net positive charge are called cations and those that acquire a net negative charge are called anions.

Oppositely charged ions attract each other to form an ionic bond, which is a chemical bond that forms arrays(indiscreet molecules) such as crystals. An example is when sodium forms an ionic bond with chloride to create sodium chloride (or table salt). An ionic bond is shown in FIGURE 2-3. Sodium has an atomic number of 11 and only has one electron inside its valence shell (its outermost energy level that contains electrons). If this electron is lost, the second shell (with eight elec-trons) becomes the valence shell. The loss of the sin-gle electron in the third (outer) shell causes sodium to become stable, meaning it is then a cation. Oppositely, chlorine (with an atomic number of 17) needs only one electron to fill its valence shell. When it gains one electron, it becomes stable and is then an anion.

The interaction of sodium and chlorine involves sodium donating one electron to chlorine. Oppositely charged ions then attract each other to form sodium chloride. Examples of ionic bonds include atoms that have one or two valence shell electrons and atoms that have seven valence shell electrons. Those with one or two valence shell electrons include calcium, potassium, and sodium (all are metallic elements) and those with seven valence shell electrons include chlorine,­ fluorine, and iodine.

In this category, most ionic compounds are referred to as salts. They do not exist as individual molecules in their dry state, but instead form crystals. These are large collections of cations and anions held together by ionic bonds. Sodium chloride gives us an example of a compound that is very different from the atoms that make it up individually. Separately, sodium is a metal that is silver white in color. Chlorine is a green-colored gas. Mixed together to become sodium chloride, the result is a white crystalline solid (table salt).

Covalent Bonds

For atoms to achieve stability, electrons do not have to be completely transferred. They may be shared, mean-ing each atom can fill its outer electron shell for part of the time. When electrons are shared, this produces molecules in which the shared electrons are located in a single orbital that is common to both atoms, which makes up a covalent bond. Hydrogen has only one electron. It can fill its only shell, labeled as Shell 1, when a pair of electrons from another atom is shared. If a hydrogen atom shares with another hydrogen atom, a molecule of hydrogen gas is formed. Therefore, the shared electron pair orbits around the molecule as awhole to make each atom achieve stability. An exampleof a covalent bond is when two hydrogen atoms bond to form a hydrogen molecule (FIGURE 2-4).

Each atom has different needs in terms of bond-ing to achieve stability. Shared electrons orbit and become part of the whole molecule, which ensures the stability of each atom. Hydrogen has only its one electron but needs two. Carbon has four electrons in its outer shell but needs eight for stability. For a methane molecule (CH4), carbon shares four pairs of electrons with four hydrogen atoms, meaning one pair with each hydrogen atom. A single covalent bond is formed when two atoms share one pair of electrons. A double covalent bond occurs when two electron pairs are shared. Likewise, a triple covalent bond occurs when three electron pairs are shared. When written, the amount of covalent bonds pres-ent are signified by using single, double, or triple horizontal lines, as (using oxygen as the example):

Polar and Nonpolar Molecules

In covalent bonds, molecules may be either polar or nonpolar. Nonpolar molecules are electrically balanced. They do not have separate positive and negative poles of charge. However, this does not always occur. Covalent bonds always have a spe-cific three -dimensional shape to their molecules. The bonds are formed at definite angles. The shape helps to determine other atoms or molecules with which the original molecule can interact. However, it may also cause unequal electron pair sharing.

This creates a polar molecule. This is common in nonsymmetric­ molecules that contain atoms with different electron­-attracting abilities.

Oxygen, nitrogen, and chlorine are examples­ of electro-hungry atoms that attract electrons strongly. This is a capability of atoms known as electronegativity­. These small atoms have six orseven valence shell electrons. However, most atoms that have only one or two valence shell electrons are electropositive. Their ability to attract electrons is solow that, most often, they lose their own valence shell electrons to other atoms. Potassium and sodium each have one valence shell electron and are examples of electropositive atoms.

Whether a covalently bonded molecule will be polar or nonpolar is determined by the molecular shape and related electron-attracting ability of each atom. For example, in the carbon dioxide molecule, four electron pairs of carbon are shared with two oxy-gen atoms, meaning two pairs are shared with each oxygen. Because oxygen is extremely electronegative, it attracts the shared electrons more strongly than car-bon is able to. Even so, because the carbon dioxide molecule is symmetric and linear, the ability of one oxygen atom to pull electrons is offset by the other oxygen atom. Therefore, the shared electrons continue to orbit the entire molecule, and carbon dioxide is a nonpolar compound.

A different example exists in the water molecule, which is V-shaped or bent. On the same end of this molecule, there are two electropositive hydrogen atoms located. The most electronegative oxygen atom is at the opposite end. Therefore, oxygen can pull the shared electrons away from the two hydrogen atoms. The electron pairs are not shared equally; instead, they are closer to the oxygen for most of the time. Because of the negative charges of electrons, the oxygen end is slightly more negative and the hydrogen end slightly more positive. Water is a polar molecule with two poles of charge, which is also called a dipole.

Polar molecules are essential for chemical reac-tions in body cells. They orient themselves toward charged particles (ions, certain proteins, etc.) or other dipoles. Various molecules have different degrees of polarity. There is a gradual change from ionic to non-polar covalent bonding. Complete electron transfer is referred to as an ionic bond. Equal electron sharing is known as a nonpolar covalent bond. These are extreme compared with each other. There are various degrees of unequal electron sharing in between these two extremes. Nonpolar covalent bonds are very com-mon and involve carbon atoms that make up most structural components of the human body.

Hydrogen Bonds

When the positive hydrogen end of a polar mole-cule is attracted to the negative nitrogen or oxygen end of another polar molecule, the attraction is called hydrogen bond (FIGURE 2-5). Hydrogen atomsalready covalently linked to one electronegative atom (such as oxygen or nitrogen) are attracted by an atom requiring electrons, forming a “bridge.” These bonds are weak at body temperature. In environmen-tal extremes, they may change form, from water to ice and back again. Hydrogen bonds are important in protein and nucleic acid structure, forming between polar regions of different parts of a single, large mol-ecule. Hydrogen bonding commonly occurs between dipoles, an example of which is the water molecule. Itoccurs because the (slightly) negative oxygen atoms of a certain molecule attract the (slightly) positive hydrogen atoms of another molecule. Therefore, because of hydrogen bonding, water molecules usu-ally cling together. They form films, and this forma-tion process is referred to as surface tension. As a result, water beads into spheres when it is on a hard surface such as a countertop. This is also the reason that certain small insects are able to walk on the sur-face of a body of water.

Hydrogen bonds are too weak to bind atoms together and form molecules. However, they are important intramolecular bonds, holding differ-ent parts of one large molecule in a specific three-dimensional shape. Proteins and DNA are examples of large biologic molecules that have many hydrogen bonds helping to maintain and stabilize their struc-tures. The water molecule is one excellent example of a hydrogen bond. It consists of two hydrogen atoms and one oxygen atom. Another example occurs when two hydrogen atoms bind with two oxygen atoms, forming hydrogen peroxide.

Because the properties of compounds are usually different from the properties of their contained atoms, it is difficult to tell which atoms are contained with-out a chemical analysis (e.g., water, which is made up of hydrogen and oxygen). The numbers and types of atoms in a molecule are represented by a molec-ular formula. The molecular formula for water is H2O, signifying­ the two atoms of hydrogen and the one atom of oxygen. Structural formulas are used to ­signify how atoms are joined and arranged inside mol-ecules. ­Single bonds are represented by single lines and double bonds are represented by double lines. When structural formulas are represented in three dimensional models, different colors are used to show different types of atoms.

1. Distinguish between ionic bonds and covalent bonds.

2. Which kind of bond holds together atoms in a water molecule?

3. Describe polar and nonpolar molecules.

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