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2.1: The multiple roles of water

  • Page ID
    14920
  • Water and other small “stuff”

    “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 the Tao Te Ching by Lao Zu.

    Let’s convert this into a chemical riddle and apply it at the nanoscopic level to biochemistry!

    “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, it donates and accepts electrons, it 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 a plethora of properties which make 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 which 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 solute in solvent would be more disordered. That simple description must be adjusted to account for ordered state of solvent (a clathrate) surrounding a solute and of “holes” in solvent the 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 entropy increase 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 as either soluble or insoluble in a given solvent. This notion can be reinforced by simple liquid/liquid partitioning experiments in organic chemistry experiments in which two immiscible solvents (for example water and an ether) are used. 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)10­CO2∟H) can form multimolecular aggregates called micelles even as they exists in as a biphasic system, as shown in the following equilibria:

    C8H17OH(liq) ↔ C8H17OH(aq) ↔ C8H17OH(micelle).

    The structure below shows a micelle. Note that all nonpolar chains are buried in the middle of the spherical structure, while the polar heads are exposed to water.

    spacer

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

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

    nC8H17OH(aq) ↔1-C8H17OH(micelle).

    To solve the general problem of limited solublity of organic molecules in aqueous based life, biomolecular structures have evolved to “carry” mostly nonpolar molecules like long chain carboxylic acids (fatty acids) and cholesterol.

    Insert iCn3D of fatty acid Protein complex and LDL or HDL particle.

    The same ideas can be applied to the solubility of salts. Students will remember general solubility rules (all Gp 1 and Gp 7 salts are soluble) from introductory chemistry. 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 Ca2+ and Fe2+ ions such as CaCO3 and FeCO3 are generally insoluble (Ksp values of 1.4 x 10-8 and 3.1 x 10-11, respectively). Insoluble calcium salts (carbonates and silicates) are need for shells of xxx and skeletons of vertebrates. Yet free Ca2+ and Fe2+ ion are found in extracellular and intracellular compartments. Divalent cations like Fe2+ can be toxic at higher concentration so ways to effectively transport and sequester them have evolved.

    iCn3D structure of ferritin?

    Finally, let’s consider the solubility of gases. The ones that are the most abundant and relevant are O2 and CO2 as they are the reactants and products of oxidative respiration. The gases, although they contain oxygen atoms, 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 that bind oxygen through bound transition metal complexes containing Fe2+ (hemoglobin in vertebrates) and Cu+ ions in hemocyanins in some invertebrates).

    Water interactions with other molecules initially begin as noncovalent interactions mediated through intermolecular forces. All biological interactions are also so mediated (except those initiated by light photons). The noncovalent interactions are weak but if they are many they can lead to very strong interactions.

    Here are the noncovalent interactions that mediate “intermolecular forces”.

    Intermolecular Forces – Noncovalent Interactions

    Interaction

    Type

    Example

    Distance
    Dependence

    Relative Strength

    (Kcal/Mol)

    Direction Dependence

    Ionic

    ion-ion.PNG

    1/r

    60

    nondirectional

    Ionic-dipole

    ion-dipole.PNG

    1/r2

    3-5

    directional

    H-Bonds

    Hbond.PNG

    4-6

    directional

    Dipole-dipole

    DipoleDipole.PNG

    1/r3

    0.5-1

    directional

    Induced Dipole-

    Induced Dipole

    InducedDi_InduceDi.PNG

    1/r6

    0.5

    (depend on size)

    nondirectional

    All of these interactions original 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:

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

    From this force derives all the electrostatic interactions described in the table above. The magnitude of the attractions for the interactions depend on the way charge is distributed in the attracting species..

    Water as a reactant: Acids and Bases

    H2O, with its sharable lone pairs and slightly positive Hs is both a Bronsted base and acid. Its acid base chemistry hence is among it’s most important features.

    Water, acting as a base, can react with both strong and weak acids. Examples of reactions of a strong acid (HCl) and weak acids (acetic acids and ammonium) with water are shown in the figure below.

    genacidbase.png

    Likewise, water can act as an acid, if it reacts with a stronger base. Here are a few examples.

    WaterAcid.PNG

    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 stronger acid/strong base to 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 actually exists in an aqueous system is H3O+(aq). This is an example of the leveling effect.

    Water as a reactant: nucleophile/electrophile

    In the reactions above, we characterized water as a Bronsted acid or base. 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 it 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 substation reaction in which water displaces ammonia as a ligand.

    TSMetalRxWater.png

    [Cu(NH3)4(H2O)2]2+ +4H2O ↔ Cu(H2O)6]2+ +4NH3

    - Hydration of an alkene. The reaction is catalyzed by addition of a proton from an acids (like H2SO4) so can be called an electrophilic hydration. Once protonated at the carbon which makes the most stable carbocation, water as a nucle0phile attacks the positively carbon to produce the alcohol.

    HydrationROH.PNG

    - Nucleophilic substation at an electrophilic carbonyl. This is a very common reaction. When water is the nucleophile, the reaction is also called a hydrolysis reaction. The reaction below is shown with OH- as the nucleophile instead of water for simplicity.

    carbonylchem_Intro.PNG

    Water as a reactant: Oxidizing/Reducing agent.

    Everyone knows what happens if you throw a piece of solid Na or K into water. An extremely exothermic reaction occurs which releases H2 gas which can catch fire and lead to an explosion. The reaction of Na is:

    2Na(s) + H2O → 2Na+(aq) + OH- (aq) + H2(g) .

    The oxidation number of elemental sodium is 0, while Na+ is +1, indicating that the sodium metal has been oxidized by the water which acts as an oxidizing agent.

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

    Zn(s) + 2H3O+(aq) ⟶ Zn2+ (aq) +2H2O(l) +H2(g)

    As in acid/base reactions, in a redox reaction, an oxidizing agent and a reducing agent react to form a new oxidizing and reducing agent.

    Other reactants can oxidize water to form oxygen. Consider fluorine gas for example:

    3F2+2H2O → O2+4HF

    F2 is a strong 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 reaction that is effectively the reverse of combustion:

    6CO2 (g) + 6H2O (l) → C6H12O6(s) + 6O2(g).

    This reaction obviously is endogonic and requires a large input of energy so the reaction proceeds to produce the potent oxidizing agent O2. The special oxygen evolving complex in photosynthesis is the powerful oxidant the can oxidize H2O to form the weaker oxidizing agent, O2.

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