🔮Chemical Basis of Bioengineering I Unit 6 – Acid-Base Equilibria & Buffers

Key Concepts

• Understand the fundamental principles of acid-base equilibria and buffer systems
• Differentiate between Arrhenius, Brønsted-Lowry, and Lewis acid-base theories
• Calculate pH and pOH values for strong and weak acids and bases
  • Use the negative logarithm of the hydrogen ion concentration to determine pH
  • Utilize the relationship between pH and pOH ($pH + pOH = 14$) to solve problems
• Interpret and apply equilibrium constants ($K_a$, $K_b$, $K_w$) in acid-base reactions
• Recognize the composition and behavior of buffer systems in maintaining stable pH
• Assess buffer capacity and prepare buffer solutions with desired pH and concentration
• Explore the significance of acid-base equilibria and buffers in biological systems and bioengineering applications
• Develop problem-solving strategies for acid-base equilibria and buffer-related calculations

Acid-Base Theories

• Arrhenius theory defines acids as H+ donors and bases as OH- donors in aqueous solutions
• Brønsted-Lowry theory expands the definition of acids as proton (H+) donors and bases as proton acceptors
  • Introduces the concept of conjugate acid-base pairs (HCl and Cl-, NH3 and NH4+)
• Lewis theory further generalizes acids as electron pair acceptors and bases as electron pair donors
• Identify the strengths of acids and bases based on their extent of dissociation in water
  • Strong acids and bases completely dissociate (HCl, NaOH), while weak acids and bases partially dissociate (CH3COOH, NH3)
• Recognize amphoteric substances that can act as both acids and bases depending on the environment (H2O, amino acids)
• Apply the appropriate acid-base theory to analyze reactions and predict products

pH and pOH

• pH is a logarithmic scale that measures the acidity or basicity of a solution
  • Defined as the negative logarithm of the hydrogen ion concentration: $pH = -log[H+]$
• pOH is the negative logarithm of the hydroxide ion concentration: $pOH = -log[OH-]$
• The pH scale ranges from 0 to 14, with 7 being neutral, < 7 acidic, and > 7 basic
• Calculate pH for strong acids and bases using the concentration of H+ or OH- ions
• Determine the pH of weak acids and bases using equilibrium constants and ICE tables
  • ICE tables help organize initial, change, and equilibrium concentrations in acid-base reactions
• Convert between pH and pOH using the relationship $pH + pOH = 14$
• Understand the significance of pH in biological systems (blood pH, enzyme activity)

Equilibrium Constants

• Equilibrium constants quantify the extent of a reaction at equilibrium
• $K_a$ is the acid dissociation constant, representing the strength of an acid
  • Larger $K_a$ values indicate stronger acids, while smaller $K_a$ values indicate weaker acids
• $K_b$ is the base dissociation constant, representing the strength of a base
  • Larger $K_b$ values indicate stronger bases, while smaller $K_b$ values indicate weaker bases
• $K_w$ is the water ionization constant, equal to the product of [H+] and [OH-] at 25°C ($K_w = [H+][OH-] = 1.0 \times 10^{-14}$)
• Calculate equilibrium concentrations using equilibrium constants and initial concentrations
• Relate $K_a$ and $K_b$ for conjugate acid-base pairs using the relationship $K_a \times K_b = K_w$
• Apply equilibrium constants to determine the pH of acid-base solutions

Buffer Systems

• Buffers are solutions that resist changes in pH when small amounts of acid or base are added
• Composed of a weak acid and its conjugate base, or a weak base and its conjugate acid
  • Examples include acetic acid (CH3COOH) and acetate ion (CH3COO-), or ammonia (NH3) and ammonium ion (NH4+)
• Maintain a relatively stable pH through the equilibrium between the weak acid and its conjugate base
• The Henderson-Hasselbalch equation relates pH, $K_a$, and the concentrations of the acid and its conjugate base: $pH = pK_a + log\frac{[A-]}{[HA]}$
• Calculate the pH of a buffer solution using the Henderson-Hasselbalch equation
• Understand the role of buffers in biological systems (blood, intracellular fluids)

Buffer Capacity and Preparation

• Buffer capacity is the ability of a buffer to resist changes in pH when acid or base is added
  • Depends on the concentrations of the weak acid and its conjugate base, as well as the $K_a$ of the acid
• Prepare buffer solutions with a desired pH and concentration using a weak acid and its conjugate base salt
  • Calculate the required amounts of acid and base using the Henderson-Hasselbalch equation and stoichiometry
• Determine the optimal buffer system for a specific pH range based on the $K_a$ values of weak acids
• Consider factors that affect buffer capacity, such as dilution and the addition of strong acids or bases
• Recognize the limitations of buffer systems in handling large pH changes or high concentrations of added acid or base

Applications in Bioengineering

• Buffers play a crucial role in maintaining pH homeostasis in biological systems
  • Blood buffer system (carbonic acid-bicarbonate) maintains pH between 7.35 and 7.45
  • Intracellular buffers (proteins, phosphates) regulate pH within cells
• pH control is essential for the proper functioning of enzymes, proteins, and cellular processes
• Buffers are used in the formulation of biopharmaceuticals to ensure stability and efficacy
  • Optimize pH conditions for drug solubility, absorption, and distribution
• Bioprocess engineering relies on buffers to maintain optimal pH for cell growth and product formation
  • Control pH in bioreactors, fermentation processes, and downstream purification steps
• Buffers are employed in the development of biomaterials and tissue engineering scaffolds
  • Maintain a suitable pH environment for cell adhesion, proliferation, and differentiation
• Design and select appropriate buffer systems for specific bioengineering applications based on pH requirements and compatibility with other components

Problem-Solving Strategies

• Identify the type of acid-base problem (strong acid/base, weak acid/base, buffer system)
• Determine the relevant equilibrium constants ($K_a$, $K_b$, $K_w$) and dissociation reactions
• Set up ICE tables to organize initial, change, and equilibrium concentrations
  • Use equilibrium constants and stoichiometry to solve for unknown concentrations
• Apply the Henderson-Hasselbalch equation for buffer system problems
  • Calculate pH, $pK_a$, or concentration ratios based on given information
• Use logarithmic properties to convert between pH, pOH, [H+], and [OH-]
• Check the reasonableness of answers by considering the pH range and the strength of acids and bases involved
• Practice solving a variety of acid-base equilibria and buffer problems to reinforce concepts and problem-solving skills
• Analyze the impact of assumptions and approximations made in problem-solving, such as neglecting the autoionization of water in certain cases