Biochemistry Lecture Notes - Water, Acids, and Bases

This set of flashcards covers key vocabulary and concepts from a biochemistry lecture focused on the properties of water, the behavior of acids and bases, and their interactions in biological systems.

Learning Outcome: Water's Solvent Properties - Polarity and Hydration

Water is described as a highly polar solvent. Its electrical distribution allows it to form hydration spheres around ions and form hydrogen bonds with polar molecules.

• Exam Note: The formation of hydration spheres is what physically separates and stabilizes ions in solution, preventing them from re-associating.

Learning Outcome: Understanding the Dielectric Constant (D)

The Dielectric Constant (D) measures a solvent's capacity to reduce the electrostatic attraction between two charges. Water has one of the highest dielectric constants (\approx 80).

• Exam Note: A high D value indicates that the solvent is very effective at shielding charges from one another, which is why ionic bonds are weaker in water than in nonpolar environments.

Learning Outcome: Application of Coulomb’s Law in Biochemistry

Coulomb's Law quantifies the energy (E) of an electrostatic interaction: E = \frac{kq1q2}{Dr}.

• q1, q2: Magnitudes of the two charges.

• D: Dielectric constant of the medium.

• r: Distance between the charges.

• Exam Note: Note that energy is inversely proportional to both the distance and the dielectric constant. Increasing either will decrease the strength of the interaction.

Learning Outcome: Characteristics of Electrostatic and van der Waals Interactions

1. Electrostatic Interactions: Includes ionic bonds (attraction between full charges) and hydrogen bonds (attraction between partial charges).

2. van der Waals Forces: Weak interactions occurring between all neutral atoms/molecules due to 'flickering' dipoles and electron fluctuations.

• Exam Note: Individually, these are 'weak' interactions, but collectively, they provide the necessary stability for large biomolecules like DNA or proteins.

Learning Outcome: Mechanism of the Hydrophobic Effect

The hydrophobic effect refers to the aggregation of nonpolar molecules when placed in water. It is not an 'attraction' between nonpolar molecules, but rather driven by water's tendency to maximize its own hydrogen bonding.

• Exam Note: This effect is the primary driving force for the folding of proteins and the assembly of the lipid bilayer in cell membranes.

Learning Outcome: Defining Acid Strength (Ka and pKa)

The Acid Dissociation Constant (K_a) measures the equilibrium of an acid's dissociation.

• K_a: A larger value means a stronger acid.

• pKa: Calculated as - \log Ka. A smaller/lower value means a stronger acid.

• Exam Note: For the exam, remember that strong acids have high Ka and low pKa, while weak acids have low Ka and high pKa.

Learning Outcome: The Logarithmic pH Scale

pH is defined as the negative logarithm of the hydrogen ion concentration: pH = -\log[H^+].

• Acidity: Lower pH values indicate higher acidity (higher concentration of H^+).

• Exam Note: Because the scale is logarithmic, a solution at pH = 5 has 100 times more H^+ than a solution at pH = 7.

Learning Outcome: Function and Composition of Biological Buffers

A buffer is a system designed to resist changes in pH when small amounts of acid (H^+) or base (OH^-) are added. It consists of a weak acid (HA) and its conjugate base (A^-).

• Exam Note: Buffers are most effective at the pH point where the concentration of the acid and conjugate base are equal (pH = pK_a).

Learning Outcome: Mathematical Relationship via Henderson-Hasselbalch

The Henderson-Hasselbalch equation relates pH to the ratio of acid and base components: pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right).

• [A^-]: Conjugate base concentration.

• [HA]: Weak acid concentration.

• Exam Note: This equation is used to calculate the resulting pH of a buffer or to determine the ratio of components needed to achieve a specific target pH.

Water as a Polar Solvent

Water is a highly polar molecule due to its electrical distribution. This polarity allows it to form hydration spheres around ions and establish hydrogen bonds with other polar molecules. These hydration spheres physically separate and stabilize ions in solution, which prevents them from re-associating into a solid.

The Dielectric Constant (D)

The Dielectric Constant measures a solvent's ability to shield and reduce the electrostatic attraction between two charges. Water has a very high dielectric constant (approximately 80). This high value explains why ionic bonds are much weaker in water than in nonpolar environments, as the water effectively shields the charges from one another.

Coulomb\'s Law in Biochemistry

Coulomb\'s Law is used to calculate the energy (E) of an electrostatic interaction: E = \frac{kq1q2}{Dr}. In this formula, q1 and q2 represent the magnitudes of the two charges, D is the dielectric constant of the medium, and r is the distance between the charges. The energy is inversely proportional to both distance and the dielectric constant.

Electrostatic vs. van der Waals Interactions

1. Electrostatic Interactions: These include ionic bonds (between full charges) and hydrogen bonds (between partial charges). 2. van der Waals Forces: These are weak, short-range interactions that occur between all neutral atoms or molecules caused by temporary \'flickering\' dipoles and electron fluctuations. Individually weak, these interactions collectively stabilize large structures like DNA.

The Hydrophobic Effect

The hydrophobic effect is the aggregation of nonpolar molecules in an aqueous environment. Rather than a direct attraction between nonpolar molecules, it is driven by water\'s need to maximize its own hydrogen bonding and entropy. This effect is the fundamental driving force behind the folding of proteins and the formation of lipid bilayers in cell membranes.

Measuring Acid Strength: Ka and pKa

Acid strength is determined by the Acid Dissociation Constant (Ka). A larger Ka indicates a stronger acid that dissociates more completely. Reversely, pKa is calculated as - \log Ka. Therefore, a smaller or lower pKa value signifies a stronger acid, while a higher pKa indicates a weaker acid.

The pH Scale and Acidity

pH represents the negative logarithm of the hydrogen ion concentration: pH = -\log[H^+]. Because the scale is logarithmic, each unit change represents a 10-fold change in acidity. For example, a solution with a pH = 5 is 100 times more acidic (10^2) than a solution with a pH = 7.

Biological Buffers: Function and Composition

A buffer is a system that resists changes in pH when small amounts of acid or base are added. It is typically composed of a weak acid (HA) and its conjugate base (A^-). Buffers are most effective at the specific pH point where the concentrations of the acid and base are equal, which occurs when pH = pK*a.

The Henderson-Hasselbalch Equation

This equation defines the relationship between pH, pKa, and the ratio of the buffer components: pH = pKa + \log\left(\frac{[A^-]}{[HA]}\right). Here, [A^-] is the concentration of the conjugate base and [HA] is the concentration of the weak acid. It is used to calculate the pH of a buffer or to prepare a buffer at a specific target pH.

Detailed Example: Dielectric Constant in Action

Consider table salt (NaCl). In the air (low D), the attraction between Na^+ and Cl^- is very strong, keeping them in a solid crystal. When placed in water (high D \approx 80), the electrostatic force between the ions is reduced by a factor of 80, allowing the hydration spheres to keep the ions separated and dissolved.

Detailed Example: The Hydrophobic Effect in Cooking

If you drop oil into water, the oil droplets merge to form one large globule. This happens because the water molecules 'push' the nonpolar oil together so the water can maintain its cohesive hydrogen-bonded network with the least amount of surface area interruption. In biology, this same 'pushing' force folds proteins into their functional shapes.

Detailed Example: pH Calculation Comparison

If solution A has a [H^+] of 10^{-4} M and solution B has a [H^+] of 10^{-6} M: 1. Solution A has a pH = 4. 2. Solution B has a pH = 6. Solution A is not just twice as acidic as B; it is 10^2 = 100 times more acidic due to the logarithmic nature of the pH scale.