There is no life without water. In this chapter, water will be used to review some very basic ideas in chemistry, particularly as applies to cell and molecular biology. What is water? H2O. Two hydrogen atoms and one oxygen atom (Figure 1). Together they form a molecule of water. They are defined as a molecule by the presence of strong chemical bonds connecting each atom.
Figure 1. (A) Oxygen and hydrogen (B) water
In this case, each atom is connected to another by a covalent bond. These are the strongest type of chemical bonds, and form when two atoms are sharing electrons in order to fill their outermost (valence) electron shell and increase stability.
The volume of an atom is defined by electrons in a very fast and energetic obit around a nucleus. The electrons are very small negatively charged particles, and the nucleus is composed of neutrons (electrically neutral) and protons (positively charged), both relatively massive in comparison to electrons. The electrons’ orbits around the nucleus can be approximated by “shells” or levels. These shells characteristically have limitations on the number of electrons that can fit within them: the first shell (closest to nucleus) holds only 2 electrons, while the second shell holds 8, and the third shell holds 18. The atom is most stable when its outer shell (and by extension, all inner ones also) is filled. The energy of the electrons also varies by level - innermost electrons have the least energy while the outermost electrons have the most.
In the case shown here, hydrogen (H) has only one electron, and for maximal stability of that electron shell, it should have two. Oxygen, on the other hand, has six electrons in its outer shell, and a filled shell would have eight. Thus, it would “like” to pull in two more electrons for maximal stability. As shown in Figure 1B, both of those requirements are fulfilled when each of the hydrogen atoms shares an electron with the oxygen, which also shares an electron each with the hydrogen. The water molecule can also be written as H—O—H, in which the single solid line indicates a pair of shared electrons, i.e. a single covalent bond. The energy of an average single covalent bond is about 80 kcal/mol. However, as shown below, double and even triple covalent bonds are possible. The strength of those types of bonds is slightly less than double (~150 kcal/mol) or triple (~200 kcal/mol) the energy of the single bonds.
Bond energy is a measure of the strength of the bond between two covalently joined atoms, and is proportional to the bond distance, which is determined by the atomic radii. It is not the same thing as bond dissociation energy, which is the energy released in a homolytic reaction (bond is split with electrons equally distributed) taking place at absolute zero, but they are similar in being measures of bond strength.
Sharing electrons is not the only way to create bonds between atoms. Ionic bonds are created when an atom donates or receives an electron, rather than sharing one. When an atom gives up an electron, the electrical balance between the numbers of positively charged protons in its nucleus and negatively charged electrons is upset, and the overall atom now has a positive electrical charge. Similarly when an atom receives an extra electron, the balance in a neutral atom is upset, and the atom becomes negatively charged. An ionic bond is formed when one atom donates an electron to an adjacent atom, creating an ionic pair, one positively and one negatively charged. The electrical attraction between the oppositely charged atoms holds them together.
Although salts (such as NaCl) are ionic compounds, not all ionic compounds are salts. The chemical definition of a salt requires that the compound be formed by the substitution of a hydrogen ion (H+ ) in the original compound. This usually occurs in neutralization reactions, such as the neutralization of hydrochloric acid, HCl (or H+ Cl-) with sodium hydroxide (Na+ (OH)- ), which yields the salt NaCl, and water (HOH = H2O).
Ionic bonds are weaker than covalent bonds, with an average bond energy of ~5.5 kcal/mol. Both covalent and ionic bonds are thermodynamically stable in dry, room temperature conditions (25°C, 298 K, 77°F). The average energy imparted when molecules collide at this temperature is only ~0.6 kcal/mol, far less than the energy needed to break a covalent or ionic bond.
Figure 2. (A) Individually, the Na atom and the Cl atom are electrically neutral. However, they are both very reactive chemically because both need only get rid of (Na) or take in (Cl) one electron to have a full outer shell. (B) Because an electron is completely transferred, the Na becomes Na+ and Cl becomes Cl- , reflecting the new charge imbalance. Although electrically no longer neutral, the thermodynamic enhancement from filling the outer shells makes both of these ions very stable.
Covalent and ionic bonds between atoms are the only way to make molecules, which are stable collections of chemically bonded atoms. However, other attractive interactions between atoms and molecules exist, but they are significantly weaker, and can be disrupted with relatively small changes in temperature or environmental conditions. These are van der Waal’s forces. They are very short-range interactions, requiring close apposition of the two atoms. As mentioned, an individual hydrogen bond (a specific type of van der Waal’s force described below) or other van der Waal’s interaction can be easily disrupted, but these types of interactions generally occur en masse. In a sense, they are like molecular Velcro® - each individual little plastic hook and individual loop of nylon could barely hold two hairs together, but a suit of velcro can hold a person on a vertical wall (a la Late Night with David Letterman, 1984).
In the case of hydrogen bonds, these occur when there is permanent asymmetric electron sharing within a covalently bonded molecule so that the shared electrons spend more time around one nucleus (thus imparting a negative character), than the other (which is therefore somewhat positive in character) to create a permanent electrical dipole. These dipole moments can interact with oppositely charged moments on other molecules or the same molecule. Van der Waals forces also include induced (nonpermanent) dipole-dipole interactions in which a temporary shift in electron density as they orbit the nucleus forms a minute charge differential, that can induce an opposite and attractive charge differential in a very close neighboring atom. In fact, some texts define van der Waals forces exclusively as such, leaving hydrogen bonds as a separate category altogether. One of the arguments for that idea is that the bond length of the average H-bond is smaller than the sum of the van der Waal’s radii of the two atoms.
As noted above, hydrogen bonds result from severely uneven sharing of electrons that generate permanent dipoles. In biological systems, this generally means that a hydrogen is covalently bound to either an oxygen or a nitrogen atom, which are both highly electronegative atoms, strongly attracting the shared electrons away from the hydrogen. Common hydrogen-bonding pairs are OH⋅⋅⋅:O, OH⋅⋅⋅:N, NH⋅⋅⋅:N, and NH⋅⋅⋅:O. Dotted lines are a common method for depicting hydrogen bonds in printed text and diagrams.
Figure 3. The hydrogen bonding of water molecules to one another is an important determinant of the physical properties of water.
Water is a molecule that has a permanent dipole (i.e. it is a polar molecule), with the highly electronegative oxygen nucleus taking the lion’s share of the shared electrons’ time, leaving the hydrogen nuclei stripped bare down to their protons. The geometry of the water molecule (Figure 1B) makes one side of the molecule somewhat negative with two pairs of free electrons, and the opposite side positive, because the shared electrons are only rarely near the hydrogen nuclei. This gives water the ability to hydrogen bond, and is the basis for several of water’s most important qualities. The ability to form many hydrogen bonds leads to a high specific heat of water, and enables it to act as a generous heat buffer. In order to get enough molecules of water moving faster and increase the temperature of the water, the energy put into the water must first be used to break apart the hydrogen bonds without generating heat. This is unlike most other liquids, which do not link internally with H-bonding. So the water is able to absorb more heat (energy) without a phase change than many other liquids.
Another important and unique characteristic of water is that the solid phase (ice) is less dense than the liquid phase. With most other liquids, as the temperature drops, the molecules have less energy, so they move less, and they stay closer together, increasing the density. Only part of that holds true with water. Again, the ability to form hydrogen bonds is directly related to this: as the temperature is lowered, the molecules move around less, affording them more opportunities to form hydrogen bonds. However, even though they are attractive, the H-bonds also act as spacers separating the water molecules more than if they were allowed to tumble about together in a liquid without forming H-bonds.
This aspect of water chemistry is actually more important to life in a geologic sense than at the cellular level. At the cellular level, the consequence is that freezing cells causes the water in them to expand and burst, killing them at low temperatures unless the cell has chemicals that act as antifreeze and lower the freezing temperature of the cytoplasm. On the other hand, at the geological level, when a pond or lake freezes in winter, the ice is less dense than water, thus staying on top of the pond, insulating deeper layers, and helping them stay liquid and able to support life (many organisms migrate deeper down in the winter). If water became more dense as it froze, as many other molecules, ice would sink, and eventually the entire pond would be completely solid, killing off most life in it once a year!
From a chemical standpoint, the polar nature of water makes it an excellent solvent for ionic and polar molecules. As you can see in the figure, the hydrogen side of water interacts with the negatively charged chloride ion, while the oxygen side of water interacts with the positively charged sodium ion, thus easily dissolving the salt. However, the polarity of water also makes it repel nonpolar molecules or by non-polar regions of molecules. This property, known as hydrophobicity, is crucial to life, since it is the basis for the formation of the biological membranes that define a cell. In general terms, the H-bonding between water molecules is very stable. Non-polar molecules cannot participate in H-bonding, and therefore create areas of instability wherever they are touching aqueous (water-based) solutions. The resolution to this problem is for hydrophobic molecules to aggregate, thus lowering the total surface area in contact with water. In living organisms, many protein and lipid molecules are amphipathic, with some portions hydrophobic, while other parts of the molecule are hydrophilic.