Molecules and Bonds

There are more than 100 known elements on Earth, but only three—oxygen, carbon, and hydrogen—make up more than 90% of the body’s mass. These three plus eight additional elements are considered major essential elements. An additional 19 minor essential elements are required in trace amounts. A periodic table showing major and minor essential elements is located inside the back cover of the book.

Most Biomolecules Contain Carbon, Hydrogen, and Oxygen

Molecules that contain carbon are known as organic molecules, because it was once thought that they all existed in or were derived from plants and animals. Organic molecules associated with living organisms are also called biomolecules. There are four major groups of biomolecules: carbohydrates, lipids, proteins, and nucleotides. The body uses the first three groups for energy and as the building blocks of cellular components. The fourth group, the nucleotides, includes DNA, RNA, ATP, and cyclic AMP. DNA and RNA are the structural components of genetic material. ATP (adenosine triphosphate) and related molecules carry energy, while cyclic AMP (adenosine monophosphate; cAMP) and related compounds regulate metabolism.

Each group of biomolecules has a characteristic composition and molecular structure. Lipids are mostly carbon and hydrogen (Fig. 2.1). Carbohydrates are primarily carbon, hydrogen, and oxygen, in the ratio CH2O (Fig. 2.2). Proteins and nucleotides contain nitrogen in addition to carbon, hydrogen, and oxygen (Figs. 2.3 and 2.4). Two amino acids, the building blocks of proteins, also contain sulfur.

DNA Replication

Not all biomolecules are pure protein, pure carbohydrate, or pure lipid, however. Conjugated proteins are protein molecules combined with another kind of biomolecule. For example, proteins combine with lipids to form lipoproteins. Lipoproteins are found in cell membranes and in the blood, where they act as carriers for less soluble molecules, such as cholesterol.

Glycosylated molecules are molecules to which a carbohydrate has been attached. Proteins combined with carbohydrates form glycoproteins. Lipids bound to carbohydrates become glycolipids. Glycoproteins and glycolipids, like lipoproteins, are important components of cell membranes [see Chapter 3].

Many biomolecules are polymers, large molecules made up of repeating units {poly-, many + -mer, a part}. For example, glycogen and starch are both glucose polymers. They differ in the way the glucose molecules attach to each other, as you can see at the bottom of Figure 2.2.

Some combinations of elements, known as functional groups, occur repeatedly in biological molecules. The atoms in a functional group tend to move from molecule to molecule as a single unit. For example, hydroxyl groups, —OH, common in many biological molecules, are added and removed as a group rather than as single hydrogen or oxygen atoms. Amino groups, —NH2, are the signature of amino acids. The phosphate group, —H2PO4, plays a role in many important cell processes, such as energy transfer and protein regulation. Addition of a phosphate group is called phosphorylation; removal of a phosphate group is dephosphorylation.

The most common functional groups are listed in Table 2.1.

Table 2.1 Common Functional Groups

Notice that oxygen, with two electrons to share, sometimes forms a double bond with another atom.

Shorthand Bond Structure
Amino —NH2
Carboxyl (acid) —COOH
Hydroxyl —OH —O—H
Phosphate —H2PO4

Concept Check

  1. List three major essential elements found in the human body.

  2. What is the general formula of a carbohydrate?

  3. What is the chemical formula of an amino group? Of a carboxyl group?

Electrons Have Four Important Biological Roles

An atom of any element has a unique combination of protons and electrons that determines the element’s properties (Fig. 2.5).

We are particularly interested in the electrons because they play four important roles in physiology:

  1. Covalent bonds. The arrangement of electrons in the outer energy level (shell) of an atom determines an element’s ability to bind with other elements. Electrons shared between atoms form strong covalent bonds that create molecules.

  2. Ions. If an atom or molecule gains or loses one or more electrons, it acquires an electrical charge and becomes an ion. Ions are the basis for electrical signaling in the body. Ions may be single atoms, like the sodium ion Na+ and chloride ion Cl. Other ions are combinations of atoms, such as the bicarbonate ion HCO3 . Important ions of the body are listed in Table 2.2.

  3. High-energy electrons. The electrons in certain atoms can capture energy from their environment and transfer it to other atoms. This allows the energy to be used for synthesis, movement, and other life processes. The released energy may also be emitted as radiation. For example, bioluminescence in fireflies is visible light emitted by high-energy electrons returning to their normal low-energy state.

  4. Free radicals. Free radicals are unstable molecules with an unpaired electron. They are thought to contribute to aging and to the development of certain diseases, such as some cancers. Free radicals and high-energy electrons are discussed later.

Table 2.2 Important Ions of the Body

Cations Anions
Na+ Sodium Cl Chloride
K+ Potassium HCO3 Bicarbonate
Ca2+ Calcium HPO4 2– Phosphate
H+ Hydrogen SO4 2– Sulfate
Mg2+ Magnesium

The role of electrons in molecular bond formation is discussed in the next section. There are four common bond types, two strong and two weak. Covalent and ionic bonds are strong bonds because they require significant amounts of energy to make or break. Hydrogen bonds and van der Waals forces are weaker bonds that require much less energy to break. Interactions between molecules with different bond types are responsible for energy use and transfer in metabolic reactions as well as a variety of other reversible interactions.

Covalent Bonds between Atoms Create Molecules

Molecules form when atoms share pairs of electrons, one electron from each atom, to create covalent bonds. These strong bonds require the input of energy to break them apart. It is possible to predict how many covalent bonds an atom can form by knowing how many unpaired electrons are in its outer shell, because an atom is most stable when all of its electrons are paired (Fig. 2.6).

For example, a hydrogen atom has one unpaired electron and one empty electron place in its outer shell. Because hydrogen has only one electron to share, it always forms one covalent bond, represented by a single line (—) between atoms. Oxygen has six electrons in an outer shell that can hold eight. That means oxygen can form two covalent bonds and fill its outer shell with electrons. If adjacent atoms share two pairs of electrons rather than just one pair, a double bond, represented by a double line (=), results. If two atoms share three pairs of electrons, they form a triple bond.

Polar and Nonpolar Molecules

Some molecules develop regions of partial positive and negative charge when the electron pairs in their covalent bonds are not evenly shared between the linked atoms. When electrons are shared unevenly, the atom(s) with the stronger attraction for electrons develops a slight negative charge (indicated by δ), and the atom(s) with the weaker attraction for electrons develops a slight positive charge (δ+). These molecules are called polar molecules because they can be said to have positive and negative ends, or poles. Certain elements, particularly nitrogen and oxygen, have a strong attraction for electrons and are often found in polar molecules.

A good example of a polar molecule is water (H2O). The larger and stronger oxygen atom pulls the hydrogen electrons toward itself. This pull leaves the two hydrogen atoms of the molecule with a partial positive charge, and the single oxygen atom with a partial negative charge from the unevenly shared electrons (Fig. 2.6b). Note that the net charge for the entire water molecule is zero. The polarity of water makes it a good solvent, and all life as we know it is based on watery, or aqueous, solutions.

A nonpolar molecule is one whose shared electrons are distributed so evenly that there are no regions of partial positive or negative charge. For example, molecules composed mostly of carbon and hydrogen, such as the fatty acid shown in Figure 2.6a, tend to be nonpolar. This is because carbon does not attract electrons as strongly as oxygen does. As a result, the carbons and hydrogens share electrons evenly, and the molecule has no regions of partial charge.

Noncovalent Bonds Facilitate Reversible Interactions

Ionic bonds, hydrogen bonds, and van der Waals forces are noncovalent bonds. They play important roles in many physiological processes, including pH, molecular shape, and the reversible binding of molecules to each other.

Ionic Bonds

Ions form when one atom has such a strong attraction for electrons that it pulls one or more electrons completely away from another atom. For example, a chlorine atom needs only one electron to fill the last of eight places in its outer shell, so it pulls an electron from a sodium atom, which has only one weakly held electron in its outer shell (Fig. 2.6c). The atom that gains electrons acquires one negative charge (−1) for each electron added, so the chlorine atom becomes a chloride ion Cl. Negatively charged ions are called anions.

An atom that gives up electrons has one positive charge (+1) for each electron lost. For example, the sodium atom becomes a sodium ion Na+. Positively charged ions are called cations.

Ionic bonds, also known as electrostatic attractions, result from the attraction between ions with opposite charges. (Remember the basic principle of electricity that says that opposite charges attract and like charges repel.) In a crystal of table salt, the solid form of ionized NaCl, ionic bonds between alternating Na+ and Cl ions hold the ions in a neatly ordered structure.

Hydrogen Bonds

A hydrogen bond is a weak attractive force between a hydrogen atom and a nearby oxygen, nitrogen, or fluorine atom. No electrons are gained, lost, or shared in a hydrogen bond. Instead, the oppositely charged regions in polar molecules are attracted to each other. Hydrogen bonds may occur between atoms in neighboring molecules or between atoms in different parts of the same molecule. For example, one water molecule may hydrogen-bond with as many as four other water molecules. As a result, the molecules line up with their neighbors in a somewhat ordered fashion (Fig. 2.6d).

Hydrogen bonding between molecules is responsible for the surface tension of water. Surface tension is the attractive force between water molecules that causes water to form spherical droplets when falling or to bead up when spilled onto a nonabsorbent surface (Fig. 2.6d). The high cohesiveness {cohaesus, to cling together} of water is due to hydrogen bonding and makes it difficult to stretch or deform, as you may have noticed in trying to pick up a wet glass that is “stuck” to a slick table top by a thin film of water. The surface tension of water influences lung function [described in Chapter 17].

Van der Waals Forces

Van der Waals forces are weak, nonspecific attractions between the nucleus of any atom and the electrons of nearby atoms. Two atoms that are weakly attracted to each other by van der Waals forces move closer together until they are so close that their electrons begin to repel one another. Consequently, van der Waals forces allow atoms to pack closely together and occupy a minimum amount of space. A single van der Waals attraction between atoms is very weak.

Concept Check

  1. Are electrons in an atom or molecule most stable when they are paired or unpaired?

  2. When an atom of an element gains or loses one or more electrons, it is called a(n)            of that element.

  3. Match each type of bond with its description:

    • (a) covalent bond

    • (b) ionic bond

    • (c) hydrogen bond

    • (d) van der
      Waals force

    1. weak attractive force between hydrogen and oxygen or nitrogen

    2. formed when two atoms share one or more pairs of electrons

    3. weak attractive force between atoms

    4. formed when one atom loses one or more electrons to a second atom