2.2: The multiple roles of water

Search Fundamentals of Biochemistry

 

Learning Goals 

(Learning goals written by Claude, Sonnet 4.6, Anthropic)

 

“Nothing in the world is as soft and yielding as it,
Yet nothing can better overcome the hard and strong,
For they can neither control nor do away with it.
The soft overcomes the hard,
The yielding overcomes the strong;”

These words come from Lao Tzu's Tao Te Ching. Let’s turn this into a chemical riddle and apply it to biochemistry at the nanoscale!

“What it loses, it can gain,
What it donates, it can accept,
It is weak yet strong,
It strengthens yet destroys.”

What is it? The answer (one of many possible) is water! It gains and loses protons, donates and accepts electrons, can be both a weaker or stronger acid or base, and can act as both an oxidizing and reducing agent, depending on the circumstances. Water, at least on our planet, appears necessary for life. We know of no biological life form that exists without it. This molecule possesses many unique properties, making it distinct from most other liquids and well-suited for the type of life on Earth. It has contrasting and oppositional properties. Let’s investigate a few.

Water as a solvent

Solubility is a property that depends on the nature of both solute and solvent. To a first approximation, we tell students in introductory chemistry and biology courses that for a solute to dissolve in a solvent and form a solution (an example of a homogenous mixture), the sum of noncovalent interactions (intermolecular forces) between solute and solvent must be greater than those among solute molecules and those among solvent molecules.

As students advance in chemistry classes, nuance is added to that general understanding as entropic contributions to solubility must be considered. Entropy is often described as the degree of apparent disorder in the system. Given that description, changes in entropy would appear to favor the soluble state, as a solution of the solute in the solvent would be more disordered. That simple description must be adjusted to account for the ordered state of solvent (a clathrate) surrounding a solute and of “holes” in the solvent that accommodate larger solute molecules. Enthalpy considerations must also be considered. The description of entropy as a measure of disorder is not precise. Rather, it should be described as a measure of the number of microstates of energy or particles available within a system. An increase in entropy would arise from an increase in the number of such available microstates, which could correlate with an increase in the disorder of a system.

Students might often consider a molecule soluble or insoluble in a given solvent. This notion can be reinforced by simple liquid-liquid partitioning experiments in organic chemistry, using two immiscible solvents (for example, water and ether). Yet diethyl ether is partially soluble in water (1 g/100 mL). Nonpolar molecules with no or few bond dipoles are generally considered insoluble. Students would know that acetic acid, a two-carbon molecule, is soluble in water, but how many carbons are necessary for the molecule to become essentially insoluble? Molecules with a single polar group (-OH, CO₂H) and a long alkyl/acyl chain are best described as amphiphilic. Amphiphiles like octanol (C₈H₁₇OH) and dodecyl sulfate (CH₃(CH₂)_10­CO₂H) can form multimolecular aggregates called micelles even as they exist in a biphasic system, as shown in the following equilibria:

\[\ce{C8H17OH(liq) ↔ C8H17OH(aq) <=> C8H17OH(micelle)}. \nonumber \]

Figure \(\PageIndex{1}\) shows an interactive iCn3D model of a micelle below, which consists of 54 self-associated molecules of dodecylphosphocholine fatty acids. It has an almost "complete" separation of polar (on the surface) and nonpolar atoms (buried).

Figure \(\PageIndex{1}\): A Micelle (Copyright; author via source). Click the image for a popup or use this link: https://structure.ncbi.nlm.nih.gov/i...Y7zSYepTQu7yV7.

Note the grey lines representing the nonpolar tails are buried from the surrounding water molecules, which form H bonds with the polar head groups.

Without some limited solubility, the following reaction could not occur:

\[\ce{nC8H17OH(aq) ↔1-C8H17OH(micelle).} \nonumber \]

To solve the general problem of the limited solubility of organic molecules in aqueous-based life, biomolecular structures have evolved to “transport” mostly nonpolar molecules like long-chain carboxylic acids (fatty acids) and cholesterol in circulation. The structure of one such fatty acid and cholesterol-containing particle, nascent high-density lipoprotein (HDL), has been determined by small-angle neutron scattering. Figure \(\PageIndex{2}\) shows an interactive iCn3D model of it. The gray sticks represent the nonpolar acyl tails of the long-chain carboxylic (fatty) acids, while the polar red (oxygen) and blue (nitrogen) atoms surrounding the surface are polar groups connected to the tails. The long magenta and dark blue "helices" represent a protein that wraps around the particle and stabilizes it.

Figure \(\PageIndex{2}\): Nascent HDL (3k2s) (Copyright; author via source). Click the image for a popup or use this external link: https://structure.ncbi.nlm.nih.gov/i...fF85h11SpeYJg6

The same ideas can be applied to the solubility of salts. From introductory chemistry, students will recall general solubility rules (all Group 1 and Group 7 salts are soluble). Salts of divalent cations are less soluble because the attractive ion-ion forces within the solid crystal lattice are too strong for the compensatory ion-dipole interactions between the ion and water. Hence, salts of Ca²⁺ and Fe²⁺ ions, such as CaCO₃ and FeCO_3, are generally insoluble (K_sp values of 1.4 x 10^-8 and 3.1 x 10^-11, respectively). Insoluble calcium salts (carbonates and silicates) are needed for the shells of Foraminifera and the skeletons of vertebrates. Yet, free Ca²⁺ and Fe²⁺ ions are found in extracellular and intracellular compartments. Divalent cations, such as Fe²⁺, can be toxic at high concentrations, so methods for effectively transporting and sequestering them have evolved. Figure \(\PageIndex{3}\) shows the structure of human heavy-chain ferritin (4zjk), a protein that forms a hollow shell in which Fe²⁺ ions (along with counter ions) are stored. The model below shows a ferritin with 120 Fe²⁺ ions (spheres) inside the hollow ferritin sphere.

Figure \(\PageIndex{3}\): Human heavy-chain ferritin (4zjk) containing 120 Fe²⁺ ions encapsulated within the hollow sphere formed by the protein.

Finally, let’s consider the solubility of gases. The most abundant and relevant ones are O₂ and CO_2, as they are the reactants and products of oxidative respiration. Although the gases contain oxygen atoms, they are nonpolar and therefore have no net dipole moment. Hence, they are quite insoluble in water. However, they must be soluble enough to allow fish to extract them from water. To solve the solubility problem, evolution has produced proteins such as vertebrate hemoglobin, which bind oxygen via a transition metal complex containing the Fe²⁺-heme complex (found in vertebrate hemoglobin). Some invertebrates use Cu ions in hemocyanins for the same purpose. Figure \(\PageIndex{4}\) shows an interactive iCn3D model of dioxygen (red spheres), bound to a planar heme (yellow highlights), which contains an Fe²⁺ at its center (not shown) in human hemoglobin (6BB5)

Figure \(\PageIndex{4}\): Oxy-Heme in Human Hemoglobin (6BB5). (Copyright; author via source). Click the image for a popup or use this external link: https://structure.ncbi.nlm.nih.gov/i...DQ5syZ2XTZ4kR9

Water engages in noncovalent interactions with itself and other molecules. Individual noncovalent interactions are weak, but when combined, they can lead to very strong interactions. You've studied noncovalent interactions before, which may have been described as “intermolecular forces.” We prefer the term noncovalent interaction. These include ion-ion, ion-dipole, hydrogen bonds, dipole-dipole, induced dipole-induced dipole, and other variants.

All of these interactions originate in the electrostatic force between two charged objects. There is only one law that describes the forces of attraction, and that’s Coulomb’s Law:

\[F=\dfrac{k Q_{1} Q_{2}}{r^{2}} \nonumber \]

From this force, all the electrostatic interactions listed above are derived. The magnitude of the attractions for these electrostatic interactions depends on the way the charge is distributed in the attracting species. We will explore these in depth in Chapter 2.4.

Water as a Reactant: Acids and Bases

H₂O, with its shareable lone pairs and slightly positive Hs, is both a Brønsted–Lowry base and acid. Its acid-base chemistry is, hence, among its most important features.

Water, acting as a base, can react with strong and weak acids. Examples of reactions of a strong acid (\(\ce{HCl}\)) and weak acids (acetic acid and ammonium) with water as a base are shown in Figure \(\PageIndex{5}\).

Figure \(\PageIndex{5}\): Reaction of acids with water as a base

Likewise, water can act as an acid as demonstrated in Figure \(\PageIndex{6}\).

Figure \(\PageIndex{6}\): Reaction of bases with water as an acid

In the first example, no net changes occur. In the second, a negatively charged deprotonated amine (a stronger base than water) can accept a proton from water, which acts as an acid. All acid/base reactions go predominantly in the direction of a stronger acid/strong base to a weaker acid/weaker base. Whether water reacts with a strong acid, such as HCl, or a weak one, like acetic acid, the strongest acid that can exist in an aqueous system is H₃O⁺_(aq). This is an example of the leveling effect.

Water as a reactant: nucleophile/electrophile

We characterized water as a Brønsted–Lowry acid or base in the reactions above. More generally, we could have said that water is a Lewis acid (an electron-pair acceptor) or a Lewis base (an electron-pair donor). In many reactions, we can also refer to water as a nucleophile (when it shares its lone pair) or an electrophile (when its slightly positive hydrogen atoms react with a nucleophile). Here are some examples.

Reaction of water with a transition metal complex.

This reaction below is effectively a nucleophilic substitution reaction in which water displaces ammonia as a ligand, as shown in Figure \(\PageIndex{7}\) and the following chemical equation.

Figure \(\PageIndex{7}\): Reaction of water with transition metal complexes

\[\ce{[Cu(NH3)4(H2O)2]^{2+} + 4H2O <=> [Cu(H2O)6]^{2+} + 4NH3 } \nonumber \]

Hydration of an alkene

The reaction is catalyzed by adding a proton from an acid (such as H₂SO₄), which can be described as an electrophilic hydration. Once protonated at the carbon, which makes the most stable carbocation, water, as a nucleophile, attacks the positive carbon to produce the alcohol. These steps are illustrated in Figure \(\PageIndex{8}\).

Figure \(\PageIndex{8}\): Mechanism for hydration of an alkene

Nucleophilic substitution at an electrophilic carbonyl

This is a very common reaction. When water acts as the nucleophile, the reaction is called hydrolysis. The reactions in Figure \(\PageIndex{9}\) are shown with OH^- as the nucleophile instead of water for simplicity.

Figure \(\PageIndex{9}\): Reaction of water (OH^-) as a nucleophile

Water as a reactant: Oxidizing/Reducing agent

Everyone knows what happens if you throw a piece of solid Na or K into water. This extremely exothermic reaction releases \(\ce{H2}\) gas, which can catch fire and lead to an explosion. The reaction of Na is:

\[\ce{2Na(s) + H2O → 2Na^{+}(aq) + OH^{-} (aq) + H2(g) .} \nonumber \]

The oxidation number of elemental sodium is 0, while Na⁺ is +1. This indicates that water oxidizes sodium metal, acting as an oxidizing agent.

This reaction occurs with many pure metals, but some that are less reactive (remember the activity series from introductory chemistry?) require acid, a protonated form of water, as shown in the reaction below:

\[\ce{Zn(s) + 2H3O^{+}(aq) ⟶ Zn^{2+} (aq) +2H2O(l) +H2(g)}\nonumber \]

As in acid-base reactions, in a redox reaction an oxidizing agent and a reducing agent form new oxidizing and reducing agents. Other reactants can oxidize water to form oxygen. Consider fluorine gas, for example:

\[\ce{3F2 + 2H2O -> O2 + 4HF}\nonumber \]

F2 is a stronger oxidizing agent (as you would surmise from its electronegativity) than O2, so the reaction proceeds vigorously to the right.

Of more biological relevance is the oxidation of water to produce O₂ in photosynthesis, a complex series of reactions that is effectively the reverse of combustion:

\[\ce{6CO2 (g) + 6H2O (l) → C6H12O6(s) + 6O2(g).}\nonumber \]

This endergonic reaction requires a large energy input and produces the potent oxidizing agent O₂. The special oxygen-evolving complex in photosynthesis is a powerful oxidant that can oxidize H₂O to form the weaker oxidizing agent O₂.

Summary

(Summary written by Claude, Sonnet 4.6, Anthropic)

This chapter reframes water not merely as a passive biological solvent but as a chemically multifunctional molecule whose contrasting and sometimes paradoxical properties make it uniquely suited to support life. Its central theme is that water's amphoteric, amphiphilic, and redox-active character pervades virtually every class of biochemical reaction.

Solubility in aqueous systems is governed by the interplay of noncovalent interactions and entropy. The familiar heuristic that "like dissolves like" is a useful starting point, but a complete treatment requires recognizing that dissolution reflects the net balance of solute-solute, solvent-solvent, and solute-solvent interactions, as well as the entropic cost of organizing solvent molecules around a solute. Molecules exist on a continuum from fully polar to fully nonpolar, and amphiphilic molecules — those bearing both polar and nonpolar regions — respond to aqueous environments by forming organized supramolecular structures such as micelles. This principle is not merely a curiosity of surface chemistry: it is the organizing logic behind lipid bilayers, lipoprotein particles, and the hydrophobic cores of folded proteins. Because many biologically essential molecules — fatty acids, cholesterol, O₂, CO₂, and divalent metal ions — are either poorly soluble or potentially toxic in free aqueous form, evolution has produced specialized protein structures (HDL, ferritin, hemoglobin, hemocyanin) to transport, store, and sequester them. All of the noncovalent interactions governing these processes — ion-ion, ion-dipole, hydrogen bonding, dipole-dipole, and London dispersion — derive from a single underlying law: Coulomb's Law.

As an acid and base, water is amphoteric in the Brønsted–Lowry sense, capable of both donating and accepting protons. All aqueous acid-base reactions proceed in the thermodynamically favored direction — from stronger acid and stronger base toward weaker acid and weaker base. The leveling effect constrains the strongest acid that can exist in water to H₃O⁺, regardless of the intrinsic acidity of the added proton source. In the Lewis sense, water functions as both an electron-pair donor (nucleophile) and acceptor (electrophile). As a nucleophile, water (or hydroxide) attacks electrophilic carbonyl carbons in hydrolysis reactions, displaces ligands from transition metal coordination spheres, and adds across activated alkenes — reactions that are mechanistically foundational to enzyme-catalyzed biochemistry. As an electrophile, water's partially positive hydrogen atoms participate in protonation steps that activate other substrates toward nucleophilic attack.

Finally, water participates directly in redox chemistry. It acts as an oxidizing agent toward reactive metals such as sodium and, under acidic conditions, toward less reactive metals. Conversely, water is oxidized to molecular oxygen during photosynthesis, a thermodynamically demanding process that requires substantial energy input and produces one of biology's most powerful oxidants. Together, these properties — solvent, acid, base, nucleophile, electrophile, oxidant, and reductant — establish water as an active chemical participant in virtually every aspect of cellular biochemistry, and as having unique physical properties resulting from an intricate balance of opposing tendencies rather than any single dominant feature.