biochem lecture 3

Chapter 2: Chemical Properties of Water

Overview of Water Chemistry

Water (H2O) is a unique and essential molecule composed of two hydrogen atoms covalently bonded to a single oxygen atom. This polar nature of water contributes to its capability to dissolve many ionic compounds and polar molecules, making it a universal solvent. Additionally, water can dissociate into hydrogen ions (H+) and hydroxide ions (OH-), which is fundamental to various chemical processes in biological systems.

Key Definitions:

Kw (Ionization Constant): At 25°C, the ion product of water is constant, represented as Kw = [H+][OH-] = 1 x 10^-14. This constant emphasizes the balance between the concentrations of H+ and OH- in pure water, contributing to the pH conceptual framework.

Acid-Base Chemistry

pH Scale:

The pH scale is a logarithmic scale that measures the acidity or basicity of a solution. The formula to calculate pH is pH = -log[H+]. In neutral pure water, the concentration of hydrogen ions is equal to that of hydroxide ions: [H+] = [OH-] = 10^-7 M, which corresponds to a pH level of 7.

Acidic and Basic Solutions:

• Acidic Solution: If the concentration of [H+] is greater than 10^-7 M, the solution exhibits a pH less than 7, indicating acidity.

• Basic Solution: Conversely, if [H+] is less than 10^-7 M, the pH exceeds 7, indicating a basic solution.

pOH:

The relationship between pH and pOH is articulated by the equation pH + pOH = 14. The pOH can also be calculated using the formula pOH = -log[OH-], providing insight into the concentration of hydroxide ions in a solution.

Properties of Various Substances

pH Values of Common Substances:

Brönsted-Lowry Acid-Base Theory

Definitions:

• Acid: A substance that donates a proton (H+).

• Base: A substance that accepts a proton.

Conjugate Pairs:

The Brönsted-Lowry theory also highlights the concept of conjugate acid-base pairs. For instance, in the reaction:HA + H2O ⇌ H3O+ + A-Here, HA is the acid, donating a proton to form its conjugate base A-.An example includes acetic acid (CH3COOH), which can donate a proton to become its conjugate base acetate (CH3COO-).

Dissociation and pKa

Dissociation Constant (Ka):

Ka indicates the strength of an acid in solution; stronger acids exhibit higher Ka values due to greater dissociation in water. The degree of dissociation of an acid correlates with its ability to release protons.

Understanding pKa:

The pKa can be calculated using the formula pKa = -log(Ka). A lower pKa signifies a stronger acid, indicating a greater tendency to donate protons. Strong acids like hydrochloric acid (HCl) completely dissociate in solution, making pKa less relevant. Conversely, weak acids, such as acetic acid, have significant pKa values that help to gauge their acidity levels.

Relationship between pH and pKa:

• Intrinsic Properties: pH is influenced by factors such as solute concentration, temperature, and pressure, which can cause shifts in the acidity of the solution.

• pKa: This is an intrinsic property of a compound and can vary particularly for different protons in polyprotic acids (compounds that can donate multiple protons).

Henderson-Hasselbalch Equation

Equation Usage:

The Henderson-Hasselbalch equation, expressed as pH = pKa + log([A-]/[HA]), serves as a tool for calculating the pH within weak acid-conjugate base systems. This equation is vital for determining acid-base balance in buffer solutions, providing a quantitative means to assess pH shifts.

Example Calculation:

For example, considering a 0.5M solution of acetic acid, if the pKa is 4.76, one can utilize the Henderson-Hasselbalch equation to determine the acid-base equilibrium based on the concentrations of the mixture components.

Buffers

Buffer Definition:

A buffer solution is characterized by its ability to resist changes in pH upon the addition of strong acids or bases. Buffers are critical in maintaining physiological pH levels within biological systems.

Buffer Components:

Effective buffers are usually composed of weak acids and bases that are near their pKa values, as they equilibrate quickly, minimizing pH fluctuations by neutralizing added acids or bases, in line with the Henderson-Hasselbalch equation.

Titration Curves:

Titration curves graphically represent the buffering capacity of solutions, with the midpoint illustrating where the buffer resists pH changes most efficiently, demonstrating its practical applications in laboratory settings.

Charge States and Amino Acids

Amino Acid Behavior:

The charge state of an amino acid is highly dependent on the surrounding pH:

• pH < pKa: The acidic form predominates (HA), indicating a positive charge.

• pH = pKa: This condition leads to equal concentrations of the acidic and basic forms, showcasing a neutral charge.

• pH > pKa: Here, the basic form is prevalent (A-), indicating a negative charge. The behavior of amino acids shifts remarkably, as seen with acetic acid (CH3COOH) shifting its form according to environmental pH.

Conclusion

Continued exploration of these chemical properties, alongside dedicated practice solving related problems, is paramount for not just understanding these concepts but applying them effectively—especially in calculating pH, pKa values, buffer systems, and weak acid titrations. This deeper comprehension fosters competency in various applications, including physiological processes and chemical reactions in laboratory contexts.