Chapter 2 - Water
Must understand water and its properties.
Why? Macromolecular components (i.e. proteins) assume shapes in response to water.
Most metabolic machinery operates in an aqueous environment.
Properties of Water
1) polarity
Covalent bonds (electron pair is shared) between oxygen and hydrogen atoms with a
bond angle of 104.5
o
.
Oxygen atom is more electronegative that hydrogen atom --> electrons spend more
time around oxygen atom than hydrogen atom --> result is a POLAR covalent bond.
Creates a permanent dipole in the molecule.
Can determine relative solubility of molecules “like dissolves like”.
2) hydrogen bonds
Due to polar covalent bonds --> attraction of water molecules for each other.
Creates hydrogen bonds = attraction of one slightly positive hydrogen atom of one
water molecule and one slightly negative oxygen atom of another water molecule.
The length of the bond is about twice that of a covalent bond.
Each water molecule can form hydrogen bonds with four other water molecules.
Weaker than covalent bonds (about 25x weaker).
Hydrogen bonds give water a high melting point.
Density of water decreases as it cools --> water expands as it freezes--> ice results
from an open lattice of water molecules --> less dense, but more ordered.
Hydrogen bonds contribute to water’s high specific heat (amount of heat needed to
raise the temperature of 1 gm of a substance 1oC) - due to the fact that hydrogen
bonds must be broken to increase the kinetic energy (motion of molecules) and
temperature of a substance --> temperature fluctuation is minimal.
Water has a high heat of vaporization - large amount of heat is needed to
evaporate water because hydrogen bonds must be broken to change water from
liquid to gaseous state.
3) universal solvent
Water can interact with and dissolve other polar compounds and those that ionize
(electrolytes) because they are hydrophilic.
Do so by aligning themselves around the electrolytes to form solvation spheres -
shell of water molecules around each ion.
Solubility of organic molecules in water depends on polarity and the ability to form
hydrogen bonds with water.
Functional groups on molecules that confer solubility:
carboxylates
protonated amines
amino
hydroxyl
carbonyl
As the number of polar groups increases in a molecule, so does its solubility in water.
4) hydrophobic interactions
Nonpolar molecules are not soluble in water because water molecules interact with
each other rather than nonpolar molecules --> nonpolar molecules are excluded and
associate with each other (known as the hydrophobic effect).
Nonpolar molecules are hydrophobic.
Molecules such as detergents or surfactants are amphipathic (have both hydrophilic
and hydrophobic portions to the molecule).
Usually have a hydrophobic chain of 12 carbon atoms plus an ionic or polar end.
Soaps are alkali metal salts of long chain fatty acids - type of detergent.
e.g. sodium palmitate
e.g. sodium dodecyl sulfate (synthetic detergent)
All form micelles (spheres in which hydrophilic heads are hydrated and hydrophobic
tails face inward.
Contain 80-100 detergent molecules.
Used to trap grease and oils inside to remove them.
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5) other noncovalent interactions in biomolecules
There are four major noncovalent forces involved in the structure and function of
biomolecules:
1) hydrogen bonds
More important when they occur between and within molecules --> stabilize
structures such as proteins and nucleic acids.
2) hydrophobic interactions
Very weak.
Important in protein shape and membrane structure.
3) charge-charge interactions or electrostatic interactions (ionic bonds)
Occur between two oppositely charged particles.
Strongest noncovalent force that occurs over greater distances.
Can be weakened significantly by water molecules (can interfere with bonding).
4) van der Waals forces
Occurs between neutral atoms.
Can be attractive or repulsive ,depending upon the distance of the two atoms.
Much weaker than hydrogen bonds.
The actual distance between atoms is the distance at which maximal attraction
occurs.
Distances vary depending upon individual atoms.
6) Nucleophilic nature of water
Chemicals that are electron-rich (nucleophiles) seek electron-deficient chemicals
(electrophiles).
Nucleophiles are negatively charged or have unshared pairs of electrons --> attack
electrophiles during substitution or addition reactions.
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Examples of nucleophiles: oxygen, nitrogen, sulfur, carbon, water (weak).
Important in condensation reactions, where hydrolysis reactions are favored.
e.g. protein ------> amino acids
In the cell, these reactions actually only occur in the presence of hydrolases.
Condensation reactions usually use ATP and exclude water to make the reactions
more favorable.
7) Ionization of water
Pure water ionizes slightly can act as an acid (proton donor) or base (proton
acceptor).
2H2O ---> H3O
+
+ OH
-
, but usually written
H2O ---> H
+
+ OH
-
Equilibrium constant for water:
Keq = [H
+
][OH
-
] = 1.8 x 10
-16
M at 25
o
C
[H2O]
if [H20] is 55.5 M --> 1 liter of H2O is 1000 g
1 mole of H2O is 18 g
Can rearrange equation to the following:
1.8 x 10
-16
M(55.5 M) = [H+][OH-]
1.0 x 10
-14
M
2
= [H+][OH-]
At equilibrium, [H+] = [OH-], so
1.0 x 10
-14
M
2
= [H+]
2
1.0 x 10
-7
= [H
+
]
8- pH scale
5
pH = - log [H+], so at equilibrium
pH = -log (1.0 x 10
-7
)
= 7
pH <7 is acidic, pH > 7 is basic or alkaline
1 change in pH units equals a 10-fold change in [H
+
]
Acid Dissociation Constants of Weak Acids
A strong acid or base is one that completely dissociates in water.
e.g. HCl ---> H+ + Cl-
A weak acid or base is one that does not; some proportion of the acid or base is dissociated,
but the rest is intact.
A weak acid or base can be described by the following equation:
weak acid (H) ----> H+ + A- conjugate acid-base pair
HA
proton donor conjugate base (conjugate acid)
Each acid has a characteristic tendency to lose its proton in solution.
The stronger the acid, the greater the tendency to lose that proton.
The equilibrium constant for this reaction is defined as the acid dissociation constant or Ka.
Ka = [H+] [conjugate base or A-]
[HA]
pKa = -logKa similar to pH
The pKa is a measure of acid strength. The more strongly dissociated the acid, the lower
the pKa, the stronger the acid.
Hence,
Ka = [H
+
] [A
-
]
[HA]
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log Ka = log [H+] [A-]
[HA]
log Ka = log [H+] + log [A-]
[HA]
-log[H+] = -log Ka + log [A-] Henderson-Hasselbach equation
[HA]
H-H equation defined the pH of a solution in terms of pKa and log of conjugate base and
weak acid concentrations.
Therefore, if [A-] = [HA], then
pH = pKa + log 1
pH = pKa
The pKa values of weak acids are determined by titration. Can calculate the pH of a solution
as increasing amounts of base are added.
e.g. acetic acid titration curve
OH-
CH3COOH ---------> CH3COO- + H2O
This is the sum of two reactions that are occurring:
H2O --------> H+ + OHCH3COOH
----> CH3COO- + H+
When add OH- to solution, will combine with free H+ ---> H2O (pH rises as [H+] falls).
When this happens, CH3COOH immediately dissociates to satisfy its equilibrium constant
(law of mass action).
As add more OH-, increase ionization of CH3COOH.
At the midpoint, 1/2 of CH3COOH has been ionized and [CH3COOH] = [CH3COO-].
As you continue to add more OH-, have a greater amount of ionized form compared to weak
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acid.
Finally reach a point where all the weak acid has been ionized.
This titration is completely reversible.
This titration curve shows that a weak acid and its anion can act as a buffer at or around
the pKa.
Important in cells where pH is critical.
Can also use this principle to determine whether amino acids are charged or not at different
pHs or just physiological pH.
Can use the H-H equation to calculate pH of a solution knowing the information in Table 2.4
(pKa values) and the ratios of the second term (don’t need to know actual concentrations,
just ratio).
If [A-] > [HA], then the pH of the solution is greater than pKa of the acid.
If [A-] < [HA], then the pH of the solution is less than the pKa of the acid.
Buffers
Solutions that prevent changes in pH when bases or acids are added.
Consist of a weak acid and its conjugate base.
Work best at + 1 pH unit from pKa --> maximal buffering capacity.
Excellent example:
blood plasma-carbon dioxide- carbonic acid- bicarbonate buffer system
CO2 + H2O ----> H2CO3 -------> HCO3
- + H+
If [H+] increases (pH falls), momentary increase in [H2CO3], and equation goes to
the left.
Excess CO2 is expired (increased respiration) to re-establish equilibrium.
Occurs in hypovolemia, diabetes, and cardiac arrest.
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If [H+] falls (pH increases), H2CO3 will dissociate to release bicarbonate ion and
hydrogen ion. This results in a fall in CO2 levels in the blood. As a result, breathing
slows.
Occurs in vomiting, hyperventilation (coming at equation from left)
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