Module 1 Lecture 8 2025 2 per page

Learning Objectives

• Recognize the characteristic shapes of different types of acid-base titration curves.

• Understand the chemical reactions occurring at different points during a titration.

• Comprehend how the acid-base chemistry of an amino acid leads to the characteristic shape of its titration curve.

Titration Overview

• A titration is an experimental technique used to determine the concentration of a solution by reacting it with a solution of known concentration.

• Essential to know the reaction stoichiometry, which reflects the mole ratio between the acid (A) and base (B).

• Common applications include:

  • Determining metal ions in body fluids.

  • Measuring CO2 levels in seawater.

  • Analyzing contaminants in industrial and food products.

  • Assessing wastewater components.

Acid/Base Titration Curves

• Observing how the pH of a weak acid such as acetic acid (CH3COOH) changes when NaOH is added helps in understanding titration curves.

• The key reaction during this titration is:

  • H3O+ + OH- → 2H2O

Titration of Acetic Acid (CH3COOH) with NaOH

1. Initial Setup:

   • 50.00 mL of 0.100 mol L-1 CH3COOH mixed with 0.100 mol L-1 NaOH.

2. Initial pH:

   • pKa of CH3COOH is 4.74; thus, initial pH is about 2.89.

3. Acidic Region:

   • As NaOH is added, CH3COOH reacts with NaOH to form CH3COO- (acetate) — pH shifts due to the dissociation of CH3COOH, acting as a buffer.

   • When 25 mL of NaOH has been added, [CH3COOH] equals [CH3COO-], resulting in:

     • pH = pKa = 4.74.

4. Equivalence Point:

   • The titration reaches completion when the number of moles of NaOH added equals the number of moles of CH3COOH initially present.

   • At this point, the pH rises to approximately 8.34 due to the formation of CH3COO- and Na+ in the solution.

5. Calculation of [CH3COO-]:

   • At the equivalence point, the reaction yields 0.005 mol of CH3COO- in a total volume of 100 mL (50 mL of acid + 50 mL of base):

     • Therefore, [CH3COO-] = 0.005 mol / 0.100 L = 0.05 mol L-1.

6. Determining pH at the Equivalence Point:

   • Using the relationship between pKa and pKb for conjugate pairs:

     • pKa(CH3COOH) + pKb(CH3COO-) = 14.

     • Calculate pKb(CH3COO-) = 14 - 4.74 = 9.26.

     • Then, Kb = 10^-9.26 ≈ 5.5 × 10^-10.

7. Calculating [OH-] at Equivalence:

   • Find [OH-] using the formula derived from Kb:

     • Kb = [OH-]² / [CH3COO-]initial.

     • After substituting, [OH-] ≈ 2.2 × 10^-6 mol L-1 leading to:

     • pOH = -log(5.2 × 10^-6) ≈ 5.28, and subsequently

     • pH = 14 - 5.28 = 8.72.

8. Alkaline Region:

   • After reaching the equivalence point, pH is controlled by the excess OH- added to the solution.

Strong Acid/Strong Base Titration Curves

1. Characteristics:

   • Strong acid/strong base titration does not exhibit a buffer region. The pH remains stable and does not change significantly even upon dilution, contrasting with weak acid/base behavior.

Diprotic Acids

• Diprotic acids can donate two protons sequentially (for example, phosphoric acid, H3PO3).

Amino Acids and Titration Curves

1. Nature of Amino Acids:

   • Comprised of both acidic and basic functional groups.

   • Characterized by a central carbon atom connected to an amino group (NH2), a carboxyl group (COOH), and a variable R group.

2. Zwitterionic Form:

   • At physiological pH, neutral amino acids exist predominantly as zwitterions (carrying both positive and negative charges).

   • Under acidic conditions, both COOH and NH3+ groups can be protonated, whilrotonated.

Glycine Titration Curve Example

1. Initial pH:

   • Glycine is more acidic than a typical weak acid, as the positively charged ion can lose a proton more easily.

2. Titration Regions Explained:

   • For less than one mole of OH- added, both COOH and NH3 groups are present, making the solution mix of both.

   • At half an equivalent of OH- (pH = pKa1 = 2.35), the curve reflects this change.

   • One mole of OH- leads to the isoelectric point (pI = 6.06); at this stage, most molecules have a net charge of zero.

   • Beyond one mole of OH- added, the remaining protons on NH3+ will start to react, tilting the balance until the completion of titration with two equivalents.

Homework Assignments

• Refer to specific questions (Problems 17.51, 17.70, 17.74, 17.100, 17.103) from the book; some may need Ka from the appendix.

e under basic conditions, both COO- and NH2 groups will be dep

Learning Objectives

• Recognize the characteristic shapes of different types of acid-base titration curves across various scenarios.

• Understand the specific chemical reactions occurring at different points during a titration, including the significance of the equivalence point and the buffer region.

• Comprehend how the acid-base chemistry of an amino acid leads to the characteristic shape of its titration curve, including how different pKa values affect behavior in titrations.

Titration Overview

A titration is an essential experimental technique used to determine the concentration of a solution by carefully reacting it with a solution of known concentration, known as a titrant. The stoichiometry of the reaction is crucial as it reflects the mole ratio between the acid (A) and base (B) involved in the titration. Common applications of titrations include:

• Determining metal ions in body fluids: for example, analyzing lead levels in blood.

• Measuring CO2 levels in seawater: which is essential for studying ocean acidification.

• Analyzing contaminants in industrial and food products: to ensure safety and compliance with regulations.

• Assessing wastewater components: to determine toxicity and treatment needs, among others.

Acid/Base Titration Curves

The observation of pH change during the titration of a weak acid, such as acetic acid (CH3COOH), when NaOH is added is essential for understanding titration curves. The key reaction during this process can be summarized as follows: H3O+ + OH- → 2H2O.

Titration of Acetic Acid (CH3COOH) with NaOH

Initial Setup:

The titration setup involves mixing 50.00 mL of 0.100 mol L-1 CH3COOH with an equal concentration of NaOH (0.100 mol L-1).

Initial pH:

The pKa of acetic acid is 4.74, leading to an initial pH around 2.89, which indicates a strongly acidic environment.

Acidic Region:

As NaOH is added incrementally, CH3COOH reacts with NaOH to form the acetate ion (CH3COO-). As more base is added:

• The pH shifts strategically due to the dissociation of CH3COOH, which buffers the solution until significant quantities of NaOH are added.

• Upon reaching the addition of 25 mL of NaOH, the concentrations of CH3COOH and CH3COO- become equal, resulting in a pH equal to pKa (4.74), marking a critical point in the titration.

Equivalence Point:

The titration is considered complete when the amount of moles of NaOH equals the initial number of moles of CH3COOH. At this equivalence point, the pH rises significantly to approximately 8.34 due to the formation of CH3COO- and Na+ ions in the solution, which transforms the solution to a basic environment.

Calculation of [CH3COO-]:

At the equivalence point, the reaction produces 0.005 mol of CH3COO- in a final total volume of 100 mL (50 mL of acid + 50 mL of base):[CH3COO-] = 0.005 mol / 0.100 L = 0.05 mol L-1.

Determining pH at the Equivalence Point:

Using the relationship between the pKa and pKb for conjugate pairs at the equivalence point:

• pKa(CH3COOH) + pKb(CH3COO-) = 14

• Calculation yields pKb(CH3COO-) = 14 - 4.74 = 9.26.

• Consequently, Kb is calculated as Kb = 10^-9.26 ≈ 5.5 × 10^-10.

Calculating [OH-] at Equivalence:

To find [OH-], we use the formula derived from Kb: Kb = [OH-]² / [CH3COO-].

• Substituting gives us [OH-] ≈ 2.2 × 10^-6 mol L-1, leading to a pOH ≈ 5.28 and thus, pH = 14 - 5.28 = 8.72.

Alkaline Region:

After surpassing the equivalence point, the solution's pH is determined by the excess OH- that has been added, indicating a fully basic solution.

Strong Acid/Strong Base Titration Curves

Characteristics of strong acid and strong base titrations are as follows:

• They do not exhibit a buffer region; rather, the pH remains stable, showing minimal deviation even when diluted, unlike behaviors observed in weak acid/base titrations.

Diprotic Acids

Diprotic acids, such as phosphoric acid (H3PO3), have the capacity to donate two protons (H+) sequentially in a titration involving multiple distinct stages.

Amino Acids and Titration Curves

Nature of Amino Acids:

Amino acids are unique organic compounds that contain both acidic and basic functional groups. They are characterized by a central carbon atom that is connected to an amino group (NH2), a carboxyl group (COOH), and a variable R group that defines the amino acid's properties.

Zwitterionic Form:

At physiological pH, neutral amino acids predominantly exist as zwitterions, carrying both positive and negative charges.

• In acidic conditions, both COOH and NH3+ groups can be protonated, while in basic conditions, both COO- and NH2 groups are likely to be deprotonated.

Glycine Titration Curve Example:

• Initial pH: Glycine is more acidic than typical weak acids, as the positively charged ion can lose a proton more readily, illustrating an early acid-base interaction.

• Titration Regions Explained: For less than one mole of OH- added, both COOH and NH3 groups dominate, resulting in a buffer solution. Upon reaching half an equivalent of OH- (where pH = pKa1 = 2.35), a significant behavioral shift is observed in the curve.

• At one mole of OH- added, the isoelectric point (pI = 6.06) is reached, where most glycine molecules possess a net charge of zero. Beyond this point, additional OH- interacts with remaining protons on the NH3+ groups until the completion of the titration with two equivalents.

Homework Assignments

• Refer to specific questions (Problems 17.51, 17.70, 17.74, 17.100, 17.103) from the textbook. Some problems may require consultation of Ka values from the appendix for accurate completion.