Date post: | 23-Dec-2015 |
Category: |
Documents |
Upload: | azkha-avicena |
View: | 222 times |
Download: | 0 times |
BIOCHEMISTRY I
Midterm exam schedule (Chap 2-7)
4/24 (Fri) 8pm
No withdrawal allowed!
• Key term report submission:
3/4 Chap 2. aqueous system.
3/9. Chap 3. amino acids, peptides, and proteins
3/16. Chap 4. tertiary and quaternary structure of proteins
3/23. Chap 5. protein functions
4/1. Chap 6. enzymes
4/8. Chap 7. carbohydrates and glycobiology
# Quizes
3/23 (Mon): Chapter 2 and 3
4/6 (Mon): Chapter 4 and 5
© 2013 W. H. Freeman and Company
2 | Water and Aqueous Solutions
© 2013 W. H. Freeman and Company
CHAPTER 2 Water and Aqueous Solutions
• What kind of interactions occur between molecules
• Why water is a good medium for life
• Why nonpolar moieties aggregate in water
• How dissolved molecules alter properties of water
• How weak acids and bases behave in water
• How buffers work and why we need them
• How water participates in biochemical reactions
Learning goals:
Water is the most abundant substances in
living system, 70% or more of the weight of
most organisms
All aspects of cell structure and function are
adapted to the physical and chemical
properties of water
Hydrogen bonding gives water its unusual properties
Water has higher melting point, boiling point, and heat of
vaporization than most other common solvents
: Hydrogen bonding
FIGURE 2-1a Structure of the water molecule. (a) The dipolar
nature of the H2O molecule; the two hydrogen atoms have localized
partial positive charges (δ+) and the oxygen atom has a partial
negative charge (δ–).
Hydrogen Bonds
• Strong dipole-dipole or charge-dipole interaction that arises
between an acid (proton donor) and a base (proton acceptor)
• Typically 4–6 kJ/mol for bonds with neutral atoms,
and 6–10 kJ/mol for bonds with one charged atom
• Typically involves two electronegative atoms (frequently
nitrogen and oxygen)
• Hydrogen bonds are strongest when the bonded molecules are
oriented to maximize electrostatic interaction
• Ideally the three atoms involved are in a line
Hydrogen Bonding in Water
• Water can serve as both
– an H donor
– an H acceptor
• Up to four H-bonds per water molecule gives water its
– anomalously high boiling point
– anomalously high melting point
– unusually large surface tension
• Hydrogen bonding in water is cooperative
• Hydrogen bonds between neighboring molecules are weak
(20 kJ/mol) relative to the H–O covalent bonds (420 kJ/mol)
Ice: Water in a Solid State
• Water has many different crystal forms;
the hexagonal ice is the most common
•Hexagonal ice forms a regular lattice,
and thus has a low entropy
•Hexagonal ice contains more hydrogen bonds/water molecule
• Thus, ice has lower density than liquid water;
• and, ice floats
Water forms hydrogen bonds with polar solutes
- Water is a polar solvent
: charged or polar compounds can be readily dissolved.
- Chloroform and benzene are nonpolar solvents
- Hydrophilic compounds: dissoves easily in water.
- Hydrophobic compounds: nonpolar molecules
Importance of Hydrogen Bonds
• Source of unique properties of water
• Structure and function of proteins
• Structure and function of DNA
• Structure and function of polysaccharides
• Binding of substrates to enzymes
• Binding of hormones to receptors
• Matching of mRNA and tRNA
Hydrogen Bonds: Examples
Biological Relevance of Hydrogen Bonds
Water dissolves many salts
• High dielectric constant reduces attraction
between oppositely charged ions in salt crystal;
almost no attraction at large (> 40 nm) distances
• Strong electrostatic interactions between the
solvated ions and water molecules lower the
energy of the system
• Entropy increases as ordered crystal lattice is
dissolved
F=Q1Q2
εr2
The strength of ionic interaction (F):
Q: magnitude of charge
ε: dielectric constant
r: the distance between charged groups
For water, at 25 °C, ε = 78.5: weak ionic interaction
For benzene, ε = 4.6 : strong ionic interaction
Ionic attractions or repulsion operate only over short distances
The Hydrophobic Effect
• Refers to the association or folding of nonpolar
molecules in the aqueous solution
• Is one of the main factors behind:
– protein folding
– protein-protein association
– formation of lipid micelles
– binding of steroid hormones to their receptors
• Does not arise because of some attractive direct
force between two nonpolar molecules
Solubility of Polar and Nonpolar Solutes
Why are nonpolar molecules poorly soluble in water?
Low solubility of hydrophobic solutes can be explained by entropy
• Bulk water has little order:– high entropy
• Water near a hydrophobic solute is highly ordered:– low entropy
Low entropy is thermodynamically unfavorable, thus hydrophobic solutes have low solubility.
Water surrounding nonpolar solutes has lower entropy
Origin of the Hydrophobic Effect (1)
• Consider amphipathic lipids in water
• Lipid molecules disperse in the solution; nonpolar
tail of each lipid molecule is surrounded by ordered
water molecules
• Entropy of the system decreases
• System is now in an unfavorable state
Origin of the Hydrophobic Effect (2)
• Nonpolar portions of the amphipathic molecule aggregate so that fewer water molecules are ordered
• The released water molecules will be more random and the entropy increases
• All nonpolar groups are sequestered from water, and the released water molecules increase the entropy further
• Only polar “head groups” are exposed and make energetically favorable H-bonds
Hydrophobic effect favors ligand binding
• Binding sites in enzymes and receptors are often hydrophobic
• Such sites can bind hydrophobic substrates and ligands such as steroid hormones
• Many drugs are designed to take advantage of the hydrophobic effect
van der Waals Interactions
• van der Waals interactions have two components:
– Attractive force (London dispersion) depends on the polarizability
– Repulsive force (Steric repulsion) depends on the size of atoms
• Attraction dominates at longer distances (typically 0.4–0.7 nm)
• Repulsion dominates at very short distances
• There is a minimum energy distance (van der Waals contact distance)
Biochemical Significance of van der Waals Interactions
• Weak individually
– easily broken, reversible
• Universal
– occur between any two atoms that are near each other
• Importance
– determines steric complementarity
– stabilizes biological macromolecules (stacking in DNA)
– facilitates binding of polarizable ligands
Examples of Noncovalent Interactions
Name and briefly define four types of
noncovalent interactions that occur between
biological molecules.
(1) Hydrogen bonds: weak electrostatic attractions between one
electronegative atom (such as oxygen or nitrogen) and a
hydrogen atom covalently linked to a second electronegative
atom
(2) electrostatic interactions: relatively weak charge-charge
interactions (attractions of opposite charges, repulsions of
like charges) between two ionized groups
(3) hydrophobic interactions: the forces that tend to bring two
hydrophobic groups together, reducing the total area of the
two groups that is exposed to surrounding molecules of the
polar solvent (water)
(4) van der Waals interactions: weak interactions between the
electric dipoles that two close-spaced atoms induce in each
other.
Weak interactions are crucial to macromolecular structure and function
Hydrogen bond
Ionic interaction
Hydrophobic interaction
Van Der Waals interaction
continually forming and breaking
=> individually insignificant but cumulatively very significant
Effects of Solutes on Properties of Water
• Colligative Properties
– Boiling point, melting point, and osmolarity
– Do not depend on the nature of the solute, just the concentration: number of solute particles
• Noncolligative Properties
– Viscosity, surface tension, taste, and color
– Depend on the chemical nature of the solute
• Cytoplasm of cells are highly concentrated solutions and have high osmotic pressure
Osmotic Pressure
Osmosis => water movement across a semipermeable membrane
driven by differences in osmotic pressure.
Isotonic solution
Hypertonic solution
Hopotonic solution
Effect of Extracellular Osmolarity
Ionization of Water
• O-H bonds are polar and can dissociate heterolytically
• Products are a proton (H+) and a hydroxide ion (OH–)
• Dissociation of water is a rapid reversible process
• Most water molecules remain un-ionized, thus pure water has very low electrical conductivity (resistance: 18 M•cm)
• The equilibrium is strongly to the left
• Extent of dissociation depends on the temperature
H2O H+ + OH-
Proton Hydration
• Protons do not exist free in solution.
• They are immediately hydrated to form hydronium ons.
• A hydronium ion is a water molecule with a proton associated with
one of the non-bonding electron pairs.
• Hydronium ions are solvated by nearby water molecules.
• The covalent and hydrogen bonds are interchangeable. This
allows for an extremely fast mobility of protons in water via
“proton hopping.”
Proton Hopping
Water chain in cytochrome f, which is part of the energy-trapping machinery of
photosynthesis in chloroplasts. Proton hopping is involved in movement of protons.
Ionization of Water: Quantitative Treatment
Concentrations of participating species in an equilibrium process are not independent but are related via the equilibrium constant:
H2O H+ + OH- Keq = ————[H+]•[OH-]
[H2O]
Keq can be determined experimentally, it is 1.8•10–16 M at 25C.[H2O] can be determined from water density, it is 55.5 M.
Ionization of water is expressed by an equilibrium constant
Keq =[H+][OH-]
[H2O]
Keq =[H+][OH-]
[55.5 M]at 25°C
( 55.5 M )( Keq ) = [H+][OH-] = Kw
Kw = [H+][OH-] = 1.0 X 10-14 M2
Ion product of water, Kw
Neutral pH:
equal amount of [H+] and [OH-] as in pure water.
[H+] = [OH-] = 10-7 M
What is pH?
• pH is defined as the negative
logarithm of the hydrogen ion
concentration
• Simplifies equations
• The pH and pOH must always
add to 14
• In neutral solution, [H+] = [OH–]
and the pH is 7
• pH can be negative ([H+] = 6 M)
pH = -log[H+]
214- M101]OH][H[ wK
14]OHlog[]Hlog[ -
14pOHpH
pH scale is logarithmic: 1 unit = 10-fold
pH of Some Common Liquids
Dissociation of Weak Electrolytes: Principle
• Weak electrolytes dissociate
only partially in water.
• Extent of dissociation is
determined by the acid
dissociation constant Ka.
• We can calculate the pH if the
Ka is known. But some
algebra is needed!
CH3
O
OH
CH3
O
O
+ H2O-+ H3O
+
Keq
]OH[ 2 eqa KK
M1074.1COOH]CH[
]COOCH][H[ 5
3
-3
aK
]COOCH[
]COOHCH[][
3
3
aKH
pKa measures acidity
• pKa = –log Ka (strong acid large Ka small pKa)
Weak acids have different pKas
Imidazole pKa = 7.0
Protonated? Deprotonated?
At pH = 8.0
At pH = 6.0
pKa=6.0
Buffers are mixtures of weak acids and their anions (conjugate base)
• Buffers resist change in pH
•At pH = pKa, there is a 50:50 mixture of acid and anion forms of the compound
•Buffering capacity of acid/anion system is greatest at pH = pKa
•Buffering capacity is lost when the pH differs from pKa by more than 1 pH unit
Acetic Acid-Acetate as a Buffer System
Henderson–Hasselbalch Equation:Derivation
HA H+ + A-
HA][
]A[logppH
-
aK
HA][
]A][H[ -
aK
]A[
HA][][H
-a
+ K
A-][
HA][log-log]log[H- aK
Chapter 2: Summary
In this chapter, we learned about:
•The nature of intermolecular forces
•The properties and structure of liquid water
•The behavior of weak acids and bases in water