BackWater, Weak Interactions, and Buffers: Structure, Properties, and Biological Importance
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Module 2: Water, Weak Interactions and Buffers
Objectives
Dissect the structure-function relationship of water molecules.
Characterize the importance of water to biological systems.
Investigate the ability of water to act as a solvent.
Explore how water influences the structure of biomolecules.
Characterize the non-covalent interactions within biomolecules.
Determine the mechanisms, and importance, of buffers and pH.
Water
General Properties
Water is the most abundant molecule in living organisms and plays both passive and active roles in biochemistry.
Passive Role: The structure and function of biomolecules are influenced by their interaction with water. For example, protein folding is driven by the tendency to bury hydrophobic residues away from water.
Active Role: Water participates directly in many biochemical reactions. For instance, peptide bond formation releases a water molecule.
Example: Formation of a peptide bond between amino acids:
Matrix of Life
Water is so critical to life that its presence is a key criterion in the search for extraterrestrial life. While water does not guarantee life, it is difficult to imagine life as we know it without water. Alternative solvents such as ammonia or formamide have been considered for hypothetical life forms.
Example: Chemical structures of ammonia (NH3) and formamide (HCONH2).
Structure/Function Relationship
The structural simplicity of water makes it an ideal molecule to illustrate the structure-function relationship in biochemistry.
Electronegativity: Oxygen is more electronegative than hydrogen, resulting in a permanent dipole in the water molecule.
Partial Charges: Oxygen carries a partial negative charge (δ−), while each hydrogen carries a partial positive charge (δ+).
Dipole Effects: The dipole allows water to form electrostatic interactions and hydrogen bonds with other molecules.
Example: Water molecule structure:
Oxygen (δ−) bonded to two hydrogens (δ+)
Hydrogen Bonds
General Features
Hydrogen bonds are electrostatic interactions between an electronegative atom (such as oxygen or nitrogen) with a hydrogen covalently linked to another electronegative atom possessing a free electron pair (acceptor).
Oxygen and nitrogen are common hydrogen bonders in biomolecules.
Both can serve as hydrogen bond donors and acceptors.
Example: Hydrogen bonds in proteins and nucleic acids stabilize secondary and tertiary structures.
Strength and Geometry
Hydrogen bonds are relatively weak compared to covalent bonds, but they are crucial for molecular structure and function.
Strength: Approximately 5% the strength of a covalent bond.
Length: Hydrogen bonds are about twice as long as covalent bonds.
Geometry: The strength depends on the geometry; anti-parallel beta sheets are more stable than parallel sheets due to optimal hydrogen bonding alignment.
Example: Beta-sheet structures in proteins.
Water: Unusual Properties
Hydrogen Bonding Capacity
Each water molecule can form up to four hydrogen bonds (two as a donor, two as an acceptor), leading to unique physical properties.
In liquid water, molecules participate in dynamic clusters of hydrogen bonds.
Hydrogen bonds confer high cohesion, affecting water's behavior and properties.
Thermal Properties
The extensive hydrogen bonding in water contributes to its high heat of vaporization and specific heat capacity.
Heat of Vaporization: The energy required to convert a liquid to a gas at its boiling point.
Specific Heat Capacity: The energy required to raise the temperature of 1 gram of a substance by 1°C.
Example: Water's high specific heat helps organisms maintain stable internal temperatures.
Physical Properties Compared to Other Solvents
Water has a higher melting point, boiling point, and heat of vaporization than most common solvents.
These properties are essential for life, allowing for temperature regulation and stability in biological systems.
Density and Ice Formation
In ice, water molecules form a highly ordered arrangement with lower density than liquid water, causing ice to float.
This property is crucial for aquatic life, as it insulates bodies of water in cold climates.
Water as a Solvent
Electrostatic Interactions
Water dissolves charged solutes by forming hydration layers around ions, stabilizing them in solution.
Water's small size and permanent dipole allow it to interact with both cations and anions.
Example: Dissolution of NaCl in water.
Hydrogen Bonding with Solutes
Many solutes have functional groups capable of hydrogen bonding with water.
Water can act as both a hydrogen bond donor and acceptor, making it an ideal solvent for polar molecules.
Solubility of Molecules
Hydrophilic molecules: Polar, charged, or capable of hydrogen bonding; highly soluble in water.
Hydrophobic molecules: Non-polar; poorly soluble in water.
Amphipathic molecules: Contain both hydrophilic and hydrophobic regions (e.g., fatty acids).
Example: Limited solubility of non-polar gases like O2 and CO2 in water.
Behavior of Amphipathic Substances
Amphipathic molecules in water arrange themselves so hydrophilic regions interact with water, while hydrophobic regions cluster together.
This leads to the formation of structures such as micelles and bilayers, fundamental to cell membranes.
Hydrophobic interactions drive the folding and stabilization of biomolecular structures.
Weak Interactions in Biomolecules
Types of Non-Covalent Interactions
Non-covalent interactions are essential for the formation, stabilization, and function of biomolecules.
Hydrogen bonds
Ionic (electrostatic) interactions
Hydrophobic interactions
van der Waals interactions
Hydrogen Bonds in Biomolecules
Functional groups in biomolecules can form hydrogen bonds with water, within the same molecule (intramolecular), or with other molecules (intermolecular).
Hydrogen bonds contribute to the specificity of molecular interactions.
Ionic (Electrostatic) Interactions
Attractive or repulsive forces between charged groups.
Strength is diminished in aqueous environments due to shielding by water molecules.
Strength depends on distance and the nature of the medium.
van der Waals Forces
Short-range interactions between permanent and induced dipoles.
Maximal attraction occurs when atoms are separated by the sum of their van der Waals radii.
Abundant in the core of folded proteins.
Hydrophobic Effect and Thermodynamics
Polar groups interact with water, while non-polar regions are shielded away.
Protein folding involves clustering of non-polar side chains in the interior and polar side chains on the surface.
Association of non-polar molecules increases the entropy of water, driving the hydrophobic effect.
Example: Folding of a polypeptide chain.
Ionization of Water and pH
Ionization of Water
Water can ionize to form hydrogen ions (H+) and hydroxide ions (OH−):
Equilibrium constant:
Ion product of water:
pH Scale
pH is a logarithmic measure of hydrogen ion concentration:
A difference of 1 pH unit corresponds to a 10-fold difference in [H+].
Acids, Bases, and Buffers
Weak Acids and Bases
Strong acids/bases dissociate completely in water; weak acids/bases do not.
The extent of dissociation is quantified by the acid dissociation constant ():
pKa is the negative logarithm of :
Titration Curves and Buffering
Titration curves reveal the pKa of weak acids.
When pH = pKa, the concentrations of acid and conjugate base are equal.
Buffering region extends one pH unit on either side of the pKa.
Example: Acetic acid (pKa = 4.76) buffers best between pH 3.76 and 5.76.
Biological Buffers
Organisms maintain constant pH to preserve biomolecular structure and function.
Blood pH is regulated by the bicarbonate buffer system:
Henderson-Hasselbalch Equation
Relates pH, pKa, and the ratio of conjugate base to acid:
Allows calculation of any one variable if the other two are known.
Example Calculation:
Given 0.01 M acetic acid and 0.1 M sodium acetate (pKa = 4.76):
Given 5 mM lactic acid and 50 mM lactate, pH = 4.80:
