Download Water Enabled Biochemistry of Life Part 1: Weak Bonds Rule!

Water Enabled Biochemistry of Life
(Textbook Chapter 2)
Part 1: Weak Bonds Rule!
Lecture 3
1
Water is life!?
https://www.youtube.com/watch?v=6OMlekqnL8I
Water is the medium for life
• Life evolved in water (UV protection)
• Organisms typically contain 70–90% water
• Chemical reactions occur in aqueous milieu
• Water is a critical determinant of the structure and function of proteins, nucleic acids, and membranes
Structure of the Water Molecule
• Octet rule dictates that there are four electron pairs around an oxygen atom in water • These electrons are in four sp3 orbitals • Two of these pairs covalently link two hydrogen atoms to a central oxygen atom
• The two remaining pairs remain nonbonding (lone pairs)
• Water geometry is a distorted tetrahedron • The electronegativity of the oxygen atom induces a net dipole moment
• Because of the dipole moment, water can serve as both a hydrogen bond donor and acceptor
Physics of Noncovalent Interactions
Noncovalent interactions do not involve sharing a pair of electrons. Based on their physical origin, one can distinguish between:
• Ionic (Coulombic) Interactions
– Electrostatic interactions between permanently charged species, or between the ion and a permanent dipole
• Dipole Interactions
– Electrostatic interactions between uncharged, but polar molecules
• van der Waals Interactions
– Weak interactions between all atoms, regardless of polarity
– Attractive (dispersion) and repulsive (steric) component
• Hydrophobic Effect
– Complex phenomenon associated with the ordering of water molecules around nonpolar substances
Examples of Noncovalent Interactions
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 as a Solvent
• Water is a good solvent for charged and polar substances
– amino acids and peptides
– small alcohols
– carbohydrates
• Water is a poor solvent for nonpolar substances
– nonpolar gases
– aromatic moieties
– aliphatic chains
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
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
“I believe that as the methods of structural chemistry are further applied to
physiological problems, it will be found that the significance of the hydrogen bond
for physiology is greater than that of any other single structural feature.”
–Linus Pauling, The Nature of the Chemical Bond, 1939
Hydrogen Bonds: Examples
Biological Relevance of Hydrogen Bonds
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)
Origin of the London Dispersion Force
• Quantum mechanical origin
• Instantaneous polarization by fluctuating charge distributions
• Universal and always attractive
• Stronger in polarizable molecules
• Important only at a short range
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
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
Water Enabled Biochemistry of Life
Part 2: It’s all about pKa
Lecture 4
Review of Acids and Bases
• Brønsted‐Lowry Acids and Bases
– Acids donate a hydrogen ion
– Bases accept a hydrogen ion
– Water is amphoteric (can act as an acid or a base)
1
2
1
2
1
1
2
2
Autoionization of Water
Dominatedbyspeciesontheleft at 25 C
Proton Hopping
Water Self‐Ionization
“Autoionization in liquid water”
Geissler PL et al. Science. 2001 291:2121. [PMID:
11251111]
https://www.youtube.com/watch?v=iqINYjHyimY
Review
pH
Defined as:
10
]
At 25⁰C:
For pure water,
10
Acidic
Neutral
Basic
What is the origin of p in pH ?
Nørby JG, “The origin and the meaning of the little p in pH”
(2000), Trends in Biochemical Sciences 25: 36-37. [PMID:
10637613]
Sørensen, S. P. L. (1909) C. R. Trav. Lab. Carlsberg 8:1–168.
Oxtoby, Principles of
Modern Chemistry, 4th Ed
pH of Some Common Liquids
Dissociation of Weak Electrolytes: Principle
O
H3C
+ H2O
O
Keq
H3 C
OH
+
O-
H3O+
K a  K eq  [H 2 O]
[H  ][CH 3COO - ]
Ka 
 1.74 10 5 M
[CH 3COOH]

[H ]  Ka 
[CH3COOH]
[CH3COO ]
• 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!
Dissociation of Weak Electrolytes: Example What is the final pH of a solution when 0.1 moles of acetic acid is added to water to a final volume of 1L?
O
Ka
H3C
O
H3C
OH
0.1 – x
Ka 
+ H+
O-
x
x
[ x ][ x ]
 1.74 10 5 M
[0.1 - x]
x 2  1.74 10 6  1.74 10 5 x
x 2  1.74 10 5 x  1.74 10 6  0
x = 0.001310,
pH = 2.883
• We assume that the only source of H+ is the weak acid
• To find the [H+], a quadratic equation must be solved Dissociation of Weak Electrolytes: Simplification O
H3C
Ka
O
+ H+
H3 C
O-
OH
0.1 – x
0.1
Ka 
x
x
x
x
[ x ][ x ]
 1.74 10 5 M
[0.1]
x 2  1.74 10 6
• The equation can be
simplified if the amount of dissociated species is much less than the amount of undissociated acid
• Approximation works for sufficiently weak acids and bases • Check that x < [total acid]
x = 0.00132,
pH = 2.880
pKa measures acidity
• pKa = –log Ka (strong acid  large Ka  small pKa)
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
Weak acids have different pKas
Henderson‐Hasselbalch Equation
Henderson‐Hasselbalch Equation
HA
H O
⇌A
H O
Biological Buffer Systems
• Maintenance of intracellular pH is vital to all cells
– Enzyme‐catalyzed reactions have optimal pH
– Solubility of polar molecules depends on H‐bond donors and acceptors
– Equilibrium between CO2 gas and dissolved HCO3– depends on pH
• Buffer systems in vivo are mainly based on – phosphate, concentration in millimolar range
– bicarbonate, important for blood plasma
– histidine, efficient buffer at neutral pH • Buffer systems in vitro are often based on sulfonic acids of cyclic amines
– HEPES
– PIPES
– CHES
HO
N
N
SO3Na
Water as a Reactant or Product in Biochemistry
Water bound to proteins is essential for their function
Water bound to proteins is essential for their function
Water channels in cell membranes
https://www.youtube.com/watch?v=GSi5-y6NHjY
Water bound to proteins is essential for their function