Polar Bonds and Molecules
A bond is considered to be polar if there is a significant difference in the electronegativities of the participating atoms. The following table gives the electronegativities of atoms that are common in biochemistry. These values can be used to estimate the partial charge on atoms in molecules. For example, since the electronegativity of hydrogen is smaller than C, S, N, and O, any bonds between hydrogen and these atoms will result in a partial positive charge on the hydrogen and a negative charge on the other atom. The larger the difference in electronegativity, the larger the difference in partial charges. Atom Electronegativity H 2.20 C 2.55 S 2.58 N 3.04 O 3.44 The dipole moment, μ, is defined by the following equation: μ=∑All atomsqr where q is the charge on the atom and r is the distance to the center of mass. A ​
polar molecule will have an overall net dipole moment. It is possible for a non­polar molecule to have polar bonds. For example, carbon dioxide (O=C=O) contains two polar bonds, but the dipole moment of one bond cancels the other, leading to no net dipole and therefore a non­polar molecule. Structure of Water
1.
Oxygen has the following electronic configuration: 1s22s22p4.
2.
In water, the 2s and the three 2p orbitals form four sp3 hybrid orbitals.
3.
These orbitals are tetrahedral in their orientation, however, the ideal bond angle of 109° is distorted to 104.5°.
4.
The orbitals are populated such that two orbitals are filled and two contain one electron each.
5.
The filled orbitals cannot form bonds and are called lone pairs of electrons.
6.
The half-filled orbitals participate in the formation of a sigma bond between oxygen and hydrogen.
7.
"Bent" water molecule generates a permanent dipole moment, making water a polar solvent.
Hydrogen Bonds
1.
H-bonds are stable because of electron sharing across the bond (i.e. a weak covalent bond) and an
electrostatic attraction between:
○
Electropositive hydrogen, attached to an electronegative atom is the hydrogen bond
donor (i.e. NH)
○
Electronegative hydrogen bond acceptor (e.g. the lone pairs of oxygen in the case of
water, or C=O group of an amide)
2.
Typical length: 1.8 Å (from hydrogen to oxygen, 2.7 Å from hydrogen to nitrogen)
3.
Typical angle: 180o ± 20o
4.
Typical energy: 20 kJ/mole.
5.
Number of hydrogen bonds depend on temperature, 4/molecule at 0oC.
The 4 possible hydrogen bonds formed with a water molecule in ice. The number of hydrogen bonds formed/molecule
in liquid water is less than four, and decreases as the temperature increases. At room temperature each water molecule
forms on average approximately 3 hydrogen bonds.
Biochemical Significance of Hydrogen Bonds:
1.
In ice, the hydrogen bonds cause the formation of cavities in the ice, lowering the density of the solid.
2.
In liquid water, the hydrogen bonds persist, and are transiently formed on a time scale of ~nano
seconds, generating small short-lived clusters of "ice" in liquid water.
3.
4.
Hydrogen bonds are present over a wide temperature range.
The hydrogen bonds in water allow water to absorb heat by breaking the hydrogen bonds without a
large increase in temperature, giving water a high heat capacity.
Solvation and Solubility
Hydrophobic​
compounds do not contain polar atoms and therefore cannot interact with water via hydrogen bonding.
Consequently, solvated hydrophobic compounds cause the formation of an ordered shell of hydrogen bonded water
molecules. Removal of the solvated hydrophobic compound will release these water molecules, increasing the entropy
of the solvent. This favorable increase in the entropy of the solvent drives the hydrophobic molecules from the aqueous
phase.
Hydrophillic​
compounds contain polar atoms, such as nitrogen or oxygen. Consequently, they can form hydrogen
bonds with water. The formation of hydrogen bonds is energetically favorable, thus hydrophilic compounds readily
dissolve in water.
Ionic compounds​
are readily solvated by water. There are two factors that favor a dissolved solution of ions over the
crystalline form.
1.
Increase in entropy of the ions. A crystal is highly ordered, with low entropy. Dissolved ions are
dispersed throughout the solution, a high entropy state.
2.
3.
4.
Electrostatic shielding. The force between two charged particles is:
F=14πεoq1q2Dr2
εo=8.854 ×10−12C2/Nm2
The force depends on the distance between the two charges and the dielectric constant (D) of the
media. A high dielectric constant, such as that found in water, is important because the forces between
charges are attenuated or reduced. Making it less favorable to have an electrostatic interaction
between the positive and negative ions.
The dielectric constant is proportional to the dipole moment of the solvent, as the dipole moment increases, D,
increases, as shown in the following table. A large dipole moment means that the solvent molecules can interact
favorably with charged solute molecules.
Compound
Dielectric Constant
Dipole Moment
Water
79
1.85
Methanol
32
1.66
Benzene
2
0.00
Amphipathic​
(or amphiphilic) compounds are both polar (or
charged) and nonpolar. An example is a fatty acid, which has a
charged carboxylate (red) and a non-polar hydrocarbon chain
(yellow). These can form micelles if the nonpolar part is
sufficiently large. Micelles are aggregates of amphipathic
molecules that sequester the nonpolar part on the inside, much
like the inside of an orange.