2.2: The multiple roles of water
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Search Fundamentals of Biochemistry
Learning Goals (ChaptGPT o1, 1/25/25)
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Interpretation of Duality in Molecular Properties:
- Analyze the paradox presented in the riddle (“weak yet strong, strengthens yet destroys”) and relate it to water’s ability to donate and accept protons/electrons.
- Explain how water's seemingly contradictory behaviors underpin its role in essential biochemical reactions.
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Water as a Solvent – Thermodynamic Considerations:
- Describe how noncovalent interactions (ion-ion, ion-dipole, hydrogen bonding, etc.) govern the solubility of molecules in water.
- Evaluate the roles of enthalpy and entropy in solvation, including the formation of clathrate structures and the impact of “holes” in the solvent.
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Solubility and Amphiphilicity:
- Investigate how amphiphilic molecules (like fatty acids) self-assemble into micelles, and the biochemical significance of these aggregates in transporting nonpolar molecules in aqueous environments.
- Discuss experimental approaches, such as liquid/liquid partitioning, to illustrate the principles of solubility and phase separation.
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Water’s Acid-Base Chemistry:
- Examine water’s role as both a Brønsted–Lowry acid and base in chemical reactions, and explain the concept of the leveling effect in aqueous systems.
- Explore acid-base equilibria involving water, including examples with strong acids (e.g., HCl) and weak acids (e.g., acetic acid).
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Nucleophilic and Electrophilic Reactions Involving Water:
- Analyze water’s dual functionality as a nucleophile and electrophile in various reaction mechanisms, such as ligand substitution in transition metal complexes and electrophilic hydration of alkenes.
- Compare and contrast the roles of water and hydroxide ions (OH⁻) in hydrolysis reactions.
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Water in Redox Chemistry:
- Explore how water acts as an oxidizing or reducing agent in reactions with metals (e.g., Na, Zn) and in biological processes like photosynthesis.
- Critically assess how water’s redox behavior contributes to generating and controlling reactive oxygen species in biological systems.
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Molecular Modeling and Visualization:
- Utilize interactive molecular models (such as iCn3D) to visualize how water interacts with biomolecules, highlighting the separation of polar and nonpolar regions in structures like micelles and lipoproteins.
- Interpret structural models of protein complexes (e.g., ferritin, hemoglobin) to understand how water’s chemical properties influence the stabilization and function of macromolecular assemblies.
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Biochemical Implications of Water’s Unique Properties:
- Develop hypotheses on how water’s contrasting properties are essential for life, linking its molecular behavior to its role in processes like nutrient transport, enzyme activity, and cellular structure formation.
- Critically assess experimental data and current research that highlight the centrality of water in maintaining biochemical homeostasis.
These goals are intended to reinforce theoretical knowledge and encourage critical thinking about how fundamental chemical principles manifest in complex biological systems.
“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 convert this into a chemical riddle and apply it to biochemistry at the nanoscopic level!
“What it loses, it gains,
What it donates, it accepts,
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/base or oxidizing/reducing agent and can lead to crystal formation or dissolution, depending on circumstances. Water, at least on our planet, appears necessary for life. We know of no biological life form that exists without it. This molecule has many unique properties, making it unique compared to most other liquids and optimal 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 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 also must 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 experiments 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, CO 2 H) and a long alkyl/acyl chain are best described as amphiphilic. Amphiphiles like octanol (C 8 H 17 OH) and dodecyl sulfate (CH 3 (CH 2 ) 10 CO 2 H) can form multimolecular aggregates called micelles even as they exist in as 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).
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.
The same ideas can be applied to the solubility of salts. From introductory chemistry, students will remember general solubility rules (all Gp 1 and Gp 7 salts are soluble). Salts of divalent cations are less soluble as 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 Ca 2 + and Fe 2 + ions, such as CaCO 3 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 shells of Foraminifera and skeletons of vertebrates. Yet free Ca 2 + and Fe 2 + ions are found in extracellular and intracellular compartments. Divalent cations like Fe 2 + can be toxic at a higher concentration, so ways to effectively transport and sequester them have evolved. Figure \(\PageIndex{3}\) shows the structure of human heavy-chain ferritin (4zjk), a protein that forms a hollow shell in which is stored Fe 2 + ions (along with counter ions). The model below shows a ferritin with 120 Fe 2 + ions (spheres) inside the hollow ferritin sphere.
Finally, let’s consider the solubility of gases. The most abundant and relevant ones are O 2 and CO 2, as they are the reactants and products of oxidative respiration. Although the gases contain oxygen atoms, they are nonpolar and have no net dipole. Hence, they are quite insoluble in water. However, they must be soluble enough to allow fish to extract it from water. To solve the solubility problem, evolution has produced proteins like vertebrate hemoglobin that bind oxygen through a transition metal complex containing the Fe 2 + -heme complex (hemoglobin in vertebrates). Some invertebrates use the transition metal 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 Fe 2 + at its center (not shown) at its center in human hemoglobin (6BB5)
Water engages in noncovalent interactions with itself and other molecules. Individual noncovalent interactions are weak, but if there are many, 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 charge is distributed in the attracting species. We will explore these in depth in Chapter 2.4.
Water as a reactant: Acids and Bases
H 2 O, with its sharable lone pairs and slightly positive Hs, is both a Brønsted–Lowry base and acid. Its acid-base chemistry, hence, is 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 acids and ammonium) with water as a base are shown in Figure \(\PageIndex{5}\).
Likewise, water can act as an acid as demonstrated in Figure \(\PageIndex{6}\).
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 H 3 O + (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 generically, we could have said water is a Lewis acid (electron pair acceptor) or Lewis base (electron-pair donor). In many reactions, we can also call water a nucleophile (when it shares its lone pair) or an electrophile (when its slightly positive H 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.
\[\ce{[Cu(NH3)4(H2O)2]^{2+} + 4H2O <=> [Cu(H2O)6]^{2+} + 4NH3 } \nonumber \]
Hydration of an alkene
The reaction is catalyzed by adding a proton from an acid (like H 2 SO 4 ), which can be called 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 \(\PageIndex{8}\).
Nucleophilic substitution at an electrophilic carbonyl
This is a very common reaction. When water is the nucleophile, the reaction is also called hydrolysis. The reactions in Figure \(\PageIndex{9}\) are shown with OH - as the nucleophile instead of water for simplicity.
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 oxidized sodium metal and acted 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 a new oxidizing and reducing agent. 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 O 2 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 O 2 . The special oxygen-evolving complex in photosynthesis is a powerful oxidant that can oxidize H 2 O to form the weaker oxidizing agent O 2 .
Summary
Chapter Summary
This chapter delves into water's unique and paradoxical role in biochemistry, using a chemical riddle as an engaging entry point. The riddle—“What it loses, it gains; What it donates, it accepts; It is weak yet strong; It strengthens yet destroys”—is a metaphor for water's ability to act in multiple, seemingly opposing ways. At the heart of the discussion is the recognition that water’s versatile chemical behavior underpins many critical processes essential to life.
Key Concepts Covered:
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Dual Nature of Water:
Water’s capacity to donate and accept protons and electrons makes it an ideal medium for countless biochemical reactions. Its ability to behave as a weak or strong acid/base and as an oxidizing or reducing agent illustrates its adaptability in various chemical environments. -
Water as a Solvent:
The chapter emphasizes water’s role as a universal solvent in biological systems. It explains how the interplay of noncovalent interactions—such as hydrogen bonds, ion-dipole, and van der Waals forces—facilitates the dissolution of solutes. The discussion extends into the nuanced contributions of enthalpy and entropy, including the formation of structured solvent shells (clathrates) and the accommodation of solute molecules. -
Solubility, Amphiphilicity, and Micelle Formation:
The text explores how solubility is not a binary trait but a spectrum influenced by the nature of both solute and solvent. It introduces amphiphilic molecules and their tendency to form micelles, a key concept for understanding how hydrophobic compounds are transported in aqueous environments. Models of micelles and lipid-protein assemblies like nascent high-density lipoproteins (HDL) illustrate this point vividly. -
Acid–Base and Nucleophilic/Electrophilic Reactions:
Water’s dual role as a Brønsted–Lowry acid/base and as a Lewis acid/base is examined through its reactions with both strong and weak acids and its participation in nucleophilic substitutions and electrophilic additions. The leveling effect in aqueous solutions is highlighted as a fundamental principle governing these reactions. -
Redox Chemistry of Water:
The chapter outlines how water acts as an oxidizing and reducing agent in reactions with metals, a particularly important feature in the context of biological redox processes. Examples include the reaction of water with alkali and transition metals and its critical role in photosynthetic water oxidation. -
Molecular Modeling and Structural Insights:
To reinforce theoretical concepts, the chapter utilizes interactive molecular models, such as those of micelles, ferritin, and hemoglobin. These visualizations help illustrate how water’s chemical properties affect the structure and function of biomolecules, influencing everything from protein stabilization to the transport of essential ions.
In summary, the chapter comprehensively examines water as more than just a solvent—it is a dynamic participant in the chemical processes that sustain life. The text provides junior and senior biochemistry majors with a deeper understanding of how water’s contrasting properties are harnessed in biological systems by integrating principles of thermodynamics, solubility, acid-base chemistry, redox reactions, and molecular interactions.