Chapter 2: Water- Non Covalent Bonds; Hydrogen Bonds & Hydrophobic interactions.

As we’ve already discussed these two before, here is a brief overview again.

3. Hydrogen bonds are among the strongest noncovalent forces in biological systems.

They are paradoxically strong enough to provide structural stability but weak enough to be readily broken.

In general, a hydrogen bond can form when a hydrogen atom covalently bonded to a strongly electronegative atom, such as nitrogenoxygen, or, in rare cases, sulfur.

Examples of hydrogen bonds that can form between molecules. Click on image for credit.

Hydrogen lies approximately 0.2 nm from another strongly electronegative atom that has an unshared electron pair. The total distance between the two electronegative atoms participating in a hydrogen bond is typically 0.27 to 0.30 nm. Some common examples of hydrogen bonds are shown.

All the functional groups are also capable of forming hydrogen bonds with water molecules. In order for hydrogen bonds to form between or within biological macromolecules, the donor and acceptor groups have to be shielded from water. In most cases this shielding occurs because the groups are buried in the hydrophobic interior of the macromolecule, where water can’t penetrate.

In DNA, for example, the hydrogen bonds between complementary base pairs are in the middle of the double helix as you can see.

Hydrogen bonds in DNA double helix. Click on image for credit.

4. Hydrophobic interactions: When relatively nonpolar molecule or group in aqueous solution associate with other nonpolar molecules rather than with water, it is termed a hydrophobic interaction.

Although hydrophobic interactions are sometimes called hydrophobic “bonds”, this description is incorrect. Nonpolar molecules or groups tend to group-up not because of mutual attraction but because the polar water molecules around them tend to pressure and entrap them close to each other as the water molecules form hydrogen bonds.

Hydrophobic interactions, like hydrogen bonds, are much weaker than covalent bonds. For example, the energy required to transfer a -CH2–  group from a hydrophobic to an aqueous environment is about 3kJ mole-1.

Again, all of the interactions covered here are individually weak compared to covalent bonds, but the combined effect of many such weak interactions can be significant.

Chapter 2: Water- More on hydrogen bonds, structure of ice and polarity.

The three dimensional interactions of liquid water is difficult to study due to it’s fluid state but much has been learned by examining the structure of the ice crystals.

Ice Crystal Lattice. Click on image for credit.

Four adjacent hydrogen-bonded oxygen atoms occupy the vertices of a tetrahedron (tetra = four, hedron = plane, thus four-sided figure). The average energy required to break each hydrogen bond in ice has been estimated to be 23 kJ mol-1.

The ability of water molecules in ice to form four hydrogen bonds and the strength of these hydrogen bonds give ice an unusually high melting point. This is because a large amount of energy, in the form of heat, is required to disrupt the hydrogen-bonded lattice of ice. When ice melts, most of the hydrogen bonds are retained by liquid water, but the bonds are distorted relative to those in ice, so that the structure of liquid water is more irregular.

Interesting point to note is that the density of most substances increase on freezing as molecular motions slows and tightly-packed crystals form. The density of water also increases as it cools until it reaches a maximum at 4°C (277K). Then, as the temperature drops below 4°C, water expands. This expansion is caused by formation of the more open hydrogen-bonded ice crystal in which each water molecule is hydrogen-bonded rigidly to four others. As a result, ice, with its open lattice, is less dense than liquid water, whose molecules can move enough to become more closely packed. Because ice is less dense than liquid water, ice floats and water freezes from the top down. This has important biological implications since a layer of ice on a pond insulates the living creatures below from extreme cold.

Furthermore, two additional properties of water are related to its hydrogen-bonded characteristics- its specific heat and its heat of vaporization.

The specific heat of substance is the amount of heat needed to raise the temperature of 1 gram of the substance by 1°C. A relatively large amount of heat is required to raise the temperature of water because each water molecule participates in multiple hydrogen bonds that must be broken in order for the kinetic energy of water molecules to increase. Btw, the abundance of water in the cells and tissues of all large multicellular organisms means that temperature fluctuations within cells are minimized. This feature is of critical biological importance since the rates of most biochemical reactions are sensitive to temperature.

The heat of vaporization of water is also much higher than that of many other liquids. A large amount of heat is required to convert water from liquid to gas because hydrogen bonds must be broken to permits water molecules to dissociate from one another and enter the gas phase. Because the evaporation of water absorbs so much heat, perspiration is an effective mechanism for decreasing body temperature; the sweat will absorb heat away from the body as it evaporates.

As discussed earlier, water molecules are permanent dipoles that are polar and can interact with and dissolve other polar compounds and compounds that ionize (the latter are called electrolytes). They can align themselves around the ions formed from electrolytes so that the negative oxygen atoms of the water molecules are oriented toward the cations (the positive ions) of the electrolytes and the positive hydrogen atoms are oriented towards the anions (the negative ions). Consider what happens when a crystal of table salt- NaCl (sodium chloride) dissolves in water. The polar water molecules are attracted to the charged ions in the crystal resulting in sodium and chloride ions on the surface of the crystal dissociating from one another, and the crystal beginning to dissolve.

Sodium Chloride dissociation (hydration) in water. Click on image for credit.

Each dissolved Na+ attracts the negative ends of several water molecules, whereas each dissolved Cl attracts the positive ends of several water molecules (see diagram).

The shell of water molecules that surrounds each ion is called a salvation sphere and usually contains several layers of solvent molecules. A molecule or ion surrounded by solvent molecules is said to be solvated. When the solvent is water, such molecules or ions are said to be hydrated.

Thus, any polar molecule has a tendency to become solvated by water molecules. Ionic organic compounds, such as carboxylates and protonated amines, owe their solubility to the amino, hydroxyl, and carbonyl groups. Molecules containing such groups disperse among water molecules, with their polar groups forming hydrogen bonds with water.

Of course, the number of polar groups in a molecule affects its solubility in water. Solubility also depends on the reaction of polar to nonpolar groups in molecule: for example, one-, two-, and three-carbon alcohols are miscible with water, but larger hydrocarbons with single hydroxyl groups are much less soluble in water.

Table measuring solubility of molecule in water as it's non-polar hydrocarbon chain grows. Click on image for credit.

In a large molecule, the properties of the nonpolar hydrocarbon portion of the molecule override those of the polar alcohol group and limit solubility.

Chapter 2: Water- Hydrogen Bonding.

All living cells depend absolutely on water for their existence. In most living cell, water is the most abundant molecule, accounting for 60% to 90% of the mass of the cell. The macromolecule components of cells-proteins, polysaccharides, nucleic acids, and membranes- get their characteristics shapes in response to interactions with water and much of the metabolic processes of cells has to operate in an aqueous environment because water is an essential solvent as well as a substrate for many cellular reactions.

A water molecule (H2O) is V-shaped, with an angle of 104.5° between the two covalent O-H bonds.

A water molecule. Click on image for credit.

An oxygen atom has six electrons in the outer shell, but the outer shell can potentially accommodate four pairs of electrons in four sp3 orbitals. This means that oxygen can form covalent bonds involving two different hydrogen atoms, each sharing a single electron with the oxygen atom.

An oxygen nucleus (because it contains more protons or positive charge) attracts electrons more strongly towards it than the single proton in the hydrogen nucleus. This attraction of electrons defines oxygen atoms as being more electronegative than hydrogen atoms. As a result, an uneven distribution of charge occurs within each O-H bond of the water molecule, with oxygen bearing a partial negative charge and hydrogen bearing a partial positive charge (+). This uneven distribution of charge within a bond is known as a dipole, and the bond is said to be polar.

Dipolarity of water bonds.

The polarity of a molecule depends on both the polarity of its covalent bonds and its geometry. The angled arrangement of the polar O-H bonds of water creates a permanent dipole for the molecule as a whole.

A molecule of ammonia also contains a permanent diploe. Thus, even though water and gaseous ammonia are electrically neutral, both molecules are polar. The high solubility of the polar ammonia molecules in water is facilitated by strong interactions with polar water molecules.

Ammonia in the form of ammonium ions. Click on image for credit.

The reason why there’s an extra proton (H+) with ammonia is because due to the fact that it is highly polar, it will attract a hydrogen atom of a water molecule (which it has dissolved in; aqueous solution) and form a fourth hydrogen bond with it. In water, the attraction between a slightly positive hydrogen atom of one water molecule and the slightly negative oxygen atom of another produces what is referred to as a hydrogen bond.

Water is not the only molecule capable of forming hydrogen bonds; these interactions can occur between any electronegative atom and a hydrogen atom attached to another electronegative atom. In the case of ammonia, the N (Nitrogen atom) is the slightly electronegative atom and the hydrogen atom of a water molecule forms a hydrogen bond with it, converting it to ammonium ions (NH4+).

The distance between this hydrogen atom and the other oxygen atom, is about twice the length of the covalent bond and hydrogen bonds are much weaker than typical covalent bonds. A single water molecule can form hydrogen bonds with up to four other water molecules.

Hydrogen Bonds in water molecules. Click on image for credit.

Orientation is important in hydrogen bonding and the bonding is most stable when a hydrogen atom and the two electronegative atoms associated with it (the two oxygen atoms in the case of water) are aligned or nearly in line.

Linear H-Bonds. Click on image for credit.