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BIOCHEMISTRY I
Water
Lecture 03
9:30 AM-10:45 AM
Outline of chapter
Hydrogen bonding in water
Some biologically important hydrogen bonds
Gibbs free energy and competition between entropy
and enthalpy
Some examples of polar, non-polar and amphipathic
compounds
water as solvent
amphipaths suspended in water
role of water in formation of enzyme-substrate
complexes (desolvation)
proton hopping
bond strengths
interactions in the aqueous environment
Ionization of water, weak acids, and weak bases
Henderson-Hasselbalch equation (pH, Pka and Buffer
concentration)
Titration of biological weak acids
Enzymatic activity and pH
buffering in the blood
Chapter 2: Water 4
Hydrogen bonding in waterRoughly tetrahedral geometry
Dipole moment
Oxygen is more electronegative and has two partial negative charges (-)
Each hydrogen nucleus (proton) has partial positive charge (+)
(a) Ball and stick model and (b) space filling model
(c) Two water models joined by a hydrogen bond (three lines)
Hydrogen bonds longer (~2-3x) and weaker (~15-20x) than covalent bond
Chapter 2: Water 5
--Waters capacity to form multiple hydrogen bonds with other water molecules alter its chemical properties with respect to other similarly sized molecules
Water Methanol EthanolBoiling point 373K 338K 352K
Heat of vaporization
40.7 kJ/mol 35.3 kJ/mol 38.6 kJ/molViscocity 1mPS 0.5mPS 1.1mPS
Surface tension 72dyn/cm 23dyn/cm 22dyn/cmDielectric constant 80.1 32.7 24.5
Chapter 2: Water 6
Hydrogen bonding in water
--Water at 0C (ice) forms H-bonds with 4 other water molecules (donates two protons and oxygen accepts two protons)
--Water at room temperature has an average of 3.4 H-bonds per water molecule (donates two protons and oxygen accepts 1.4 protons)
HYDROGEN BONDING IN BIOLOGICAL SYSTEMS
Hydrogen bonding is not unique to water, viz.:
In fact: H-bonds can form between an electronegative atom and an H-donor moiety. The latter is an H atom bonded to another electronegative atom, typically N, O, or F .; i.e., the
ensemble consists of H sandwiched between two electronegative atoms
Examples of H-bond structures ubiquitous in biological systems:
Directionality of the Hydrogen Bond
electrostatic interactions maximized when acceptor atom and H atom are in a straight line
directionality contributes, i.e., to the ability of proteins and nucleic acids (with many such bonds) to form precise 3D structures
Intermolecular vs Intramolecular Hydrogen bonding
Chapter 2: Water 9
Rudimentary Thermodynamics
G = H - TS
Reactants Products
Keq = [Products] / [Reactants]
G = -RT ln (Keq)
G is the Gibbs free energy:
If G = (-) then Keq is large and Products are favored
If G = (-) then Keq is small and reactants are favored
H is change in enthalpy (internal energy of system). H = (-) means that the system is more stable
S is change in entropy (disorder)
Water as a solvent
Chapter 2: Water 9
Water is a polar solvent and dissolves most biological molecules that are hydrophilic (water-loving).
Hydrophobic (water-fearing) and non-polar molecules like lipids and waxes dissolve poorly in water.
Water can dissolve salts like NaCl (below) by hydrating and stabilizing the ions. This weakens the electrostatic interactions of the ions acting against the ions tendency to form a crystalline lattice.
Water as a solvent
random orientation
Chapter 2: Water 10
Amphipathic compounds contain polar (or charged) regions and nonpolar regions.
Suspension of amphipaths in water
The hydrophobic alkyl groups cluster together to present the smallest hydrophobic area to water, and the polar head groups arrange to maximize their interaction with water.
Chapter 2: Water 11
H-bonds between water molecules are broken (H = Hwater - Hmixture = +)
Water molecules are forced to form a cage around the hydrophobic tail (S = -)
G = H - TS (H = +; S = -; thus
G = +)
This is an unfavorable reaction. Lipids have low solubility in water (biochemists/biologists refer to lipids as being suspended in water rather than dissolved in water)
Why is S = - when it is + in the NaCl example?
Chapter 2: Water 12
--Nonpolar regions cluster together to minimize contact with water (hydrophobic interactions) (H = -, but small)
--Polar regions arrange to maximize interaction with water--Driving force is water, not hydrophobic clustering/packing!
--Thermodynamic stability gained by minimizing # of ordered H2O molecules S = + and large
Chapter 2: Water 13
------------
--release of water molecules is an increase in entropy (S = +)
packing of hydrophobic chains is decrease in enthalpy (H = -)
Critical micelle concentration (CMC): The concentration at which micelles begin to readily form. A property of each lipid -think of it as Keq.
Chapter 2: Water 14
Hydrophobic molecules (e.g. hexane, testosterone) dissolve in water via a similar concept.
Water forms an ordered cage (clathrate structure) around
hydrophobic molecules.
Chapter 2: Water 15
--Enzyme and substrate have a shell of water that hydrates/surrounds them
--Binding of enzyme to substrate releases some of this. This is ann increase in entropy of water (S = +)
--Decrease in enthalpy (H = -) as a result of more H-bonds, ionic interaction, and hydrophobic interaction.
----------------------------
Water in biochemistry
Chapter 2: Water 17
Proteins arrange water molecules next to one another like chain/wire
Chain of water molecules allows for protons to hop from one end to
the other end
Proton hopping is faster than diffusion
Used by some membrane proteins to regulate influx of protons into
organelles
Hydronium ion: (H3O+)
Proton hopping is much faster than diffusion explains the high mobility of H+ ions compared to other noncovalent cations such as Na+ or K+.
Proton Hopping
Bond strengthsCovalent bonds (sharing of electron pairs between atoms. ~300 - 400 kJ/mol)
19Chapter 2: Water
Non-covalent bonds (~3-40 kJ/mol)
Interactions in the aqueous environmentNon-covalent bond interactions in water (3-40 kJ/mol)
20
Ionic EquilibriaAcids are proton donors
Bases are proton acceptors
Many biological molecules are weak acids or weak bases
!The ionization of water and the ion product
[H2O] + [H2O] [H3O+]+[OH-]
[H2O] [H+]+[OH-]
[H+][OH-] = Kw = 1 x 10-14 M2
[H+]=[OH-] = 1 x 10-7 M
The dissociation of an acid can be written as
HA+ A + H+HA A- + H+HA- A2- + H+
The equilibrium constant (Ka) is defined as
Ka = [H+][A-]/[HA]
The strength of the acid is usually expressed in terms of pKa
pKa = -log(Ka)
Chapter 2: Water 25
pH = - log [H+]
pX = - log [X]
------------------------------------------------------
-----------------------------------------------------Neutral
acidic
basic
--logarithmic, not arithmetic
--1 pH unit = 10X change in [H+]
--pH affects activity of biological macromolecules (e.g., enzymes)
--pH of blood or urine commonly used in medical diagnoses (acidosis vs. alkalosis)
Chapter 2: Water 26
pH of Common Aqueous Liquids
Acids are proton donors
Bases are proton acceptors
A proton donor and its corresponding acceptor are called a conjugate acid-base pair
The Henderson-Hasselbalch EquationDescribes the chemical composition of a buffer as a function of pH
[HA] is the concentration of the weak acid
[A-] is the concentration of the conjugate base
[H+] is the concentration of proton
Ka is a known quantity particular for an acid/base
Chapter 2: Water 27
[HA] [H+] [A-]
deeStamp
Henderson-Hasselbalch Equation
pKa is a measurement of the tendency for a weak acid (of a
conjugate acid-base pair) to lose a proton in aqueous media
if the pH is less than the pKa (by at least one pH unit), then
the H-HEqn indicates that the weak acid form (i.e., the proton
donor) will be the dominant species ([HA]>>[A]-)
if the pH=pKa, then the weak acid (proton donor) and its
conjugate base (proton acceptor) will be present in equal
amounts ([HA]=[A]-)--a.k.a.,"half-equivalence point". Also,
effective pH buffer region (i.e., pH range resistant to change
in pH upon addition of acid or base) = pKa 1
if pH>pKa (by at least one pH unit), then the conjugate base
(proton acceptor will be the dominant species ([A]->> [HA])
this is reason why at physiological pH (~7), amino acids are
zwitterionic, viz.,
Chapter 2: Water 28
Titration of weak acidsHenderson-Hasselbalch Equation: relates pH, pK, and buffer concentration. If you know any two the third can be calculated
[A-]
pH = pKa + log [proton acceptor][proton donor]
acetate ion
Acetic acid[HA]
What is the pH of a buffer mixture containing 1M acetic acid and 0.5M sodium acetate?
The Ka of acetic acid is 1.74 x 10-5 M
The blue box is the pH at which acetic acid has its greatest buffering capacity
Histidine Titration Curve
Titration of Three Weak Acids
Weakest Acid
Ammonia
Ammonium ion
Acetic acidStrongest acid (loses proton most readily)
Di-H
Mono-H
The buffering capacity of any weak acid/base is its pKa 1
A weak acid is 1% ionized at 2 pH units below its pKa and 99% ionized at 2 pH units above its pKa
Chapter 2: Water 29
Enzymes typically have a pH optimum, as seen for pepsin (active in stomach), trypsin (acts in small intestine), and alkaline phosphatase (active in bone, role in bone mineralization).
Chapter 2: Water 30
Enzymatic activity and pH
Chapter 2: Water 31
--The bicarbonate buffer system is important in blood
--Three reversible reactions: CO2 in lungs affects equilibrium
--Hyperventilation: Excessive breathing reduces CO2 in blood and increases pH of blood. Hypoventilation is the opposite!
(bicarbonat
(carbonic
(dissolved
ga
pK=6.37
Buffering the blood