G. Watson Student_Notes-2

Unit 8 - Acids & Bases Overview

8.1 Introduction to Acids and Bases

Arrhenius Definitions:

• Acids: Substances that release H+ ions (protons) when dissolved in water. For example, Hydrochloric acid (HCl) dissociates in water as follows:

  • HCl(g) → H+ (aq) + Cl− (aq)

• Bases: Substances that release OH− ions (hydroxide ions) in water. For example, Sodium hydroxide (NaOH) dissociates in water as:

  • NaOH(s) → Na+ (aq) + OH− (aq)

Strong vs. Weak Acids and Bases:

• Strong Acids: These fully ionize in solution, releasing all of their H+ ions. Common examples include:

  • Hydrochloric acid (HCl)

  • Sulfuric acid (H2SO4)

• Weak Acids: These only partially ionize in solution. The degree of ionization can be represented by the dissociation constant (Ka). A prevalent example is carbonic acid (H2CO3).

Practice Problems:

1. Identify whether the following substances are acids or bases and classify them as strong or weak:a) HNO3b) CH3COOHc) NaOHd) HF

2. Write the dissociation equation for the following acids:a) H2SO4b) H3PO4

Conceptual Problems:

• Situational Problem: You are tasked with preparing a buffer solution that can maintain a stable pH during a fermentation process. Discuss how you would select the appropriate weak acid and its conjugate base from common laboratory chemicals.

• Conceptual Understanding: Explain how the Arrhenius definition of acids and bases helps in understanding the role of acid-base reactions in biological systems, like digestion.

8.2 pH and pOH of Strong Acids and Bases

Calculations:

• pH: The negative logarithm of the hydrogen ion concentration. The formula used is:

  • pH = -log[H+]

• pOH: The negative logarithm of the hydroxide ion concentration. The formula used is:

  • pOH = -log[OH−]

• Ion Product Constant for Water (Kw): At 25°C, Kw = [H3O+][OH−] = 10^-14. The relationship between pH and pOH is expressed as:

  • pKw = pH + pOH = 14

Practice Problems:

1. Calculate the pH of a solution where the hydrogen ion concentration is [H+] = 1.00 x 10^-3 M.Answer: pH = -log(1.00 x 10^-3) = 3.00

2. If the pH of a solution is 4.00, find the pOH.Answer: pOH = 14 - 4.00 = 10.00

3. Determine the [OH−] concentration in a solution with a pH of 9.Answer: pOH = 14 - 9 = 5; [OH−] = 10^-5 M

Conceptual Problems:

• Situational Problem: A chemist monitors the pH during the titration of a strong acid with a strong base. Discuss how the pH changes at the beginning, middle, and end of the titration, and what this indicates about the reaction progress.

• Application Scenario: Consider a swimming pool with a pH of 8. How could you use the concepts of pH and pOH to adjust the water to optimal conditions?

8.3 Weak Acid and Base Equilibria

Equilibrium in Weak Acids:

• The dissociation of a weak acid can be represented as:

  • HA ⇌ H+ + A−

• The acid dissociation constant (Ka) quantifies the extent of dissociation:

  • Ka = [H+][A−]/[HA]

Equilibrium in Weak Bases:

• The dissociation of a weak base is represented as:

  • B + H2O ⇌ BH+ + OH−

• The base dissociation constant (Kb) indicates the extent of the weak base’s ionization:

  • Kb = [BH+][OH−]/[B]

Practice Problems:

1. For the weak acid acetic acid (CH3COOH), if 0.1 M acid is initially present and 0.01 M H+ is found at equilibrium, calculate Ka.Answer: Ka = [H+][A−]/[HA] = (0.01)(0.01) / (0.1 - 0.01) = 0.001/0.09 = 0.0111 M

2. Given the dissociation reaction of ammonia (NH3 + H2O ⇌ NH4+ + OH−), if [OH−] at equilibrium is 0.01 M, calculate Kb for ammonia.Answer: Kb = [NH4+][OH−]/[NH3] = (0.01)(0.01) / (0.1 - 0.01) = 0.0001/0.09 = 0.001111 M

Conceptual Problems:

• Situational Problem: A buffer containing a weak acid and its conjugate base is used to maintain a specific pH in a laboratory experiment. Discuss how changing the concentration of either component would affect the buffer’s ability to resist pH changes.

• Conceptual Understanding: Explain how Le Chatelier's principle applies to the dissociation of weak acids and the effect of common ions on their ionization.

8.4 Acid-Base Reactions and Buffers

Buffer Solutions:

• Buffers are solutions that can resist changes in pH when small amounts of acids or bases are added. They typically comprise:

  • A weak acid and its conjugate base (e.g., acetic acid and sodium acetate) or a weak base and its conjugate acid.

  • Example: HC2H3O2 + NaC2H3O2 → Buffer solution.

Practice Problems:

1. Describe how adding a strong acid would impact the pH of a buffer solution containing acetic acid and sodium acetate, and explain why the buffer resists changes in pH.Answer: Adding a strong acid lowers pH, but the buffer resists pH changes by neutralizing the added H+ with the acetate ion.

2. Calculate the pH of a buffer consisting of 0.1 M acetic acid and 0.1 M sodium acetate.Answer: pH = 4.76 + log(0.1/0.1) = 4.76

Conceptual Problems:

• Situational Problem: A biologist is studying how changes in pH can affect enzyme activity in metabolic pathways, which are buffered systems. Discuss why maintaining a stable pH is essential for enzyme function.

• Real-World Application: Discuss how blood maintains its pH through buffer systems and the consequences of failing to maintain pH within the physiological range.

8.5 Acid-Base Titrations

Endpoint vs. Equivalence Point:

• Equivalence Point: This is the point in a titration when stoichiometric amounts of acid and base have reacted, resulting in complete neutralization.

• Indicators: These are chosen based on their pKa values, which should align with the expected pH at the equivalence point.

Titration Curves:

• Graphs plotting pH against titrant volume provide insight into how pH varies with the addition of titrant.

Practice Problems:

1. During a titration of 50.0 mL of 0.1 M HCl with 0.1 M NaOH, calculate the pH at the equivalence point.Answer: At the equivalence point, pH is 7.00 (neutral solution).

2. Sketch the expected titration curve for the titration of a strong acid with a strong base and label the equivalence point.Answer: A sharp vertical rise near pH 7, equivalence point marked.(Visual graph not provided in text)

8.6 Molecular Structure of Acids and Bases

Factors Affecting Acid Strength:

• The ability of an acid to donate protons is influenced by the stability of its conjugate base; a more stable conjugate base translates to a stronger acid.

Inductive Effect:

• The presence of electronegative atoms in the vicinity can stabilize the negative charge on conjugate bases, thus increasing acidity.

Practice Problems:

1. Compare the acidity of HF, HCl, HBr, and HI based on the stability of their conjugate bases.Answer: HI > HBr > HCl > HF; stability of conjugate base increases down the group.

2. Explain how the inductive effect of fluorine atoms attached to a carboxylic acid affects its acidity.Answer: The electronegative fluorine stabilizes the conjugate base, increasing acidity.

8.7 pH and pKa

Henderson-Hasselbalch Equation:

• This equation is vital for calculating the pH of buffer solutions and is given as:

  • pH = pKa + log([A−]/[HA])

  • Here, [A−] represents the concentration of the conjugate base and [HA] is the concentration of the weak acid.

Practice Problems:

1. If a buffer solution contains 0.2 M acetic acid and 0.2 M sodium acetate, use the Henderson-Hasselbalch equation to calculate the pH of the solution.Answer: pH = 4.76 + log(0.2/0.2) = 4.76

2. A solution has a pKa of 4.76. If the concentrations of the conjugate acid and base are equal, what is the pH of the solution?Answer: pH = pKa = 4.76

8.8 Properties of Buffers

Buffer Capacity:

• Buffer capacity measures a buffer’s effectiveness in resisting pH changes and is influenced by the concentrations of its constitutive acid and base. Higher concentrations lead to increased buffer capacity, allowing it to absorb more added acid or base.

Practice Problems:

1. Discuss how the concentrations of the acid and conjugate base in a buffer affect its buffer capacity.Answer: Higher concentrations increase buffer capacity, allowing resistance to pH changes.

2. Given a buffer with 0.15 M acetic acid and 0.15 M sodium acetate, describe its capacity to withstand pH changes compared to a buffer with 0.05 M concentrations.Answer: The 0.15 M buffer has higher capacity due to greater amounts of buffer components.

8.9 Henderson-Hasselbalch Equation

• This equation is instrumental in determining the pH of buffer solutions based on the ratio of acid to conjugate base concentrations.

Practice Problems:

1. For a buffer solution containing 0.25 M acetic acid and 0.05 M acetate, calculate the pH using the Henderson-Hasselbalch equation (pKa of acetic acid is 4.76).Answer: pH = 4.76 + log(0.05/0.25) = 4.76 - 0.301 = 4.46

8.10 Buffer Capacity

• Buffers maintain pH levels effectively within narrow limits. However, their capacity diminishes significantly if the addition of acid or base alters the concentration ratios more than a certain threshold.

Practice Problems:

1. Explain the concept of buffer capacity and its importance in biological systems.Answer: Buffer capacity measures how well a buffer can maintain pH, vital for physiological functions.

2. How would adding a large amount of strong acid to a buffer affect its ability to maintain pH?Answer: It may exceed the buffer capacity, resulting in significant pH changes.

8.11 pH and Solubility

Relation Between pH and Solubility:

• The solubility of salts in solution can vary with the pH level; for example, basic anions can enhance solubility as the pH decreases, demonstrating the critical interplay between acid-base chemistry and solubility dynamics.

Practice Problems:

1. Discuss how the solubility of calcium carbonate (CaCO3) varies with pH changes in solution.Answer: Solubility increases in acidic solutions as H+ ions react with carbonate ions.

2. Given that the solubility of a salt increases in acidic solutions, explain this phenomenon in terms of Le Chatelier's principle.Answer: Adding acid shifts equilibrium to dissolve more salt, according to Le Chatelier's principle.

Cross-Sectional Practice Problems

1. A weak acid has a Ka of 1.0 x 10^-5. Calculate the pH of a 0.1 M solution of this weak acid.

   • Answer: pH ≈ 4.00

2. You start with 0.1 M acetic acid and titrate it with 0.1 M NaOH. What is the pH at the equivalence point?

   • Answer: pH ≈ 9.00

3. If the pKa of acetic acid is 4.76, what is the concentration ratio of acetate to acetic acid in a buffer with a pH of 5.26?

   • Answer: Ratio = 1:3 (using the Henderson-Hasselbalch equation)

4. Explain why a buffer is effective in a biological system.

   • Answer: Buffers stabilize pH crucial for enzymatic reactions.

5. You have a solution of 0.2 M acetic acid and 0.1 M sodium acetate. What is the pH?

   • Answer: pH = 4.76 + log(0.1/0.2) = 4.76 - 0.301 = 4.46

6. A buffer containing 0.2 M acetic acid and 0.1 M acetate has 0.01 M HCl added. What is the resulting pH?

   • Answer: pH will slightly decrease, but remain around 4.45.

7. Determine the solubility of lead(II) chloride in an acidic solution in which [H+] = 0.10 M.

   • Answer: Higher solubility due to formation of PbCl+ complex in acidic conditions.

8. Explain the impact of pH on the solubility of calcium phosphate in the body.

   • Answer: Increased solubility in slightly acidic conditions due to H+ ions reacting with phosphate.

9. During a titration of acetic acid with NaOH, how does the pH change during the addition of the base?

   • Answer: Initial gradual increase, then steep rise near equivalence point.

10. Why does the addition of a strong acid to a buffer only slightly change its pH?

• Answer: The buffer components neutralize the added acid, minimizing pH change.

11. Calculate the pKa of a weak acid if the pH of the solution is 5.00 and the concentration of its conjugate base is 0.1 M and the weak acid is 0.05 M.

• Answer: pKa = 5.00 - log(0.1/0.05) = 5.00 + 0.301 = 5.30

12. Describe the changes in pH during the titration of a strong base into a weak acid.

• Answer: pH increases steadily, with a steep rise near the equivalence point above 7.0.

13. Calculate the pOH of a 0.005 M NaOH solution.

• Answer: pOH = -log(0.005) ≈ 2.30

14. If a solution of acetic acid is partially neutralized with NaOH, how does this influence the pH?

• Answer: The pH will rise, but not above 7, due to the buffer effect.

15. Evaluate the strength of H2SO4 compared to acetic acid.

• Answer: H2SO4 is a strong acid (fully ionized), while acetic acid is weak.

16. Discuss the importance of the Henderson-Hasselbalch equation in clinical settings.

• Answer: It helps in determining the optimal pH for physiological functions and drug activity.

17. How does temperature affect the pKa of acetic acid?

• Answer: pKa decreases with an increase in temperature, affecting buffer capacity.

18. Predict how the addition of NaOH affects pH in a solution of butanoic acid (pKa = 4.82).

• Answer: The pH will increase but remain below 7 until equivalence point is reached.

19. State how a buffer can prevent pH fluctuations in gastrointestinal juices.

• Answer: Buffers maintain pH close to neutral for optimal enzyme activity.

20. Explain why a higher concentration of buffer components leads to a greater buffer capacity.

• Answer: Greater amounts provide more resistance to pH changes upon addition of acids or bases.

21. How will the pH change upon dilution of a buffer solution?

• Answer: The pH remains relatively stable upon dilution as long as the ratio of acid to conjugate base remains unchanged.

22. Describe the role of buffers in blood physiology.

• Answer: Buffers regulate blood pH, crucial for enzyme function and metabolic processes.

23. Calculate the pH of a solution formed by mixing equal volumes of 0.1 M HCl and 0.1 M NaOH.

• Answer: pH = 7.00, indicating complete neutralization.