Weak Base-Strong Acid Titrations: Videos & Practice Problems

In a weak base-strong acid titration, the weak base (analyte) reacts with a strong acid (titrant), causing a decrease in pH. The equivalence point occurs when moles of acid equal moles of base, leading to the formation of a conjugate acid, resulting in a pH less than 7. Calculating pH involves using ICF charts for moles and ICE charts for weak species. The Henderson-Hasselbalch equation can be applied for buffer solutions formed during the titration. Understanding these concepts is crucial for mastering acid-base reactions and titration techniques.

Weak Base-Strong Acid Titrations

In a weak base-strong acid titration, the weak base serves as the analyte, while the strong acid acts as the titrant. As the strong acid is added to the weak base, the pH of the solution decreases due to the introduction of protons from the acid. This titration process necessitates the use of an ICF chart, which stands for Initial, Change, and Final, to track the changes in concentration throughout the titration. Unlike an ICE chart that uses molarity, the ICF chart utilizes moles as its unit of measurement.

To effectively conduct a titration, it is crucial to determine the equivalence volume, which is the volume of titrant needed to reach the equivalence point where the moles of acid equal the moles of base. The relationship between moles, volume, and molarity is expressed by the equation:

$m^{\mathrm{acid}}Vm{\mathrm{base}}^V$

For example, in the titration of 100 mL of 0.100 M ammonia with 0.20 M hydrochloric acid, the equivalence volume can be calculated. The moles of acid and base at the equivalence point can be set equal, leading to the calculation of the volume of hydrochloric acid needed, which is found to be 50 mL.

Initially, before any strong acid is added, the solution contains only the weak base. To find the pH of this solution, an ICE chart is employed. In this case, the weak base ammonia (NH₃) reacts with water to form ammonium ions (NH₄⁺) and hydroxide ions (OH^-). The initial concentrations are set up in the ICE chart, where the change in concentration is represented as:

$-x,+x,+x$

Using the base dissociation constant (K_b), the equilibrium expression can be established as:

$K{b}_{}=x^2(\mathrm{[NH₃]}-x)$

In this scenario, if the value of K_b for ammonia is approximately 1.76 × 10^-5, the 5% approximation method can be applied to simplify calculations. By dividing the initial concentration of ammonia by K_b, if the resulting ratio exceeds 500, the term "-x" can be ignored, streamlining the process. This leads to the calculation of hydroxide ion concentration (OH^-), which can then be used to find pOH and subsequently pH:

$\mathrm{pOH}=\; -\backslash log(\backslash text\{[OH-]\})$

From the calculated pOH, the pH can be derived using the relationship:

$\mathrm{pH}=\; 14\; -\; \backslash text\{pOH\}$

In this case, the pH is determined to be approximately 11.123 before any strong acid is added. This foundational understanding sets the stage for observing how the pH changes as the strong acid is gradually introduced into the solution.

Weak Base-Strong Acid Titrations

In a titration involving a strong acid and a weak base, such as hydrochloric acid (HCl) and ammonia (NH₃), understanding the pH changes throughout the process is crucial. To reach the equivalence point, 50 mL of HCl is required to completely neutralize 100 mL of 0.100 M ammonia. Initially, when only 20 mL of HCl is added, the system is still before the equivalence point, and the pH can be determined using an ICF (Initial, Change, Final) chart.

To calculate the moles of the reactants, the formula for moles is used: moles = liters × molarity. Converting milliliters to liters by dividing by 1000 and multiplying by the molarity gives the moles of ammonia and hydrochloric acid. In this case, the moles of ammonia (the weak base) and hydrochloric acid (the strong acid) are compared. The smaller amount will be subtracted from the larger, resulting in 0.006 moles of the conjugate base remaining after the reaction, while the strong acid is completely consumed.

According to the law of conservation of mass, the total amount of matter remains constant, merely changing form during the reaction. At the end of the titration, a weak acid (the conjugate acid) and a conjugate base are present, which together form a buffer solution. The pH of this buffer can be calculated using the Henderson-Hasselbalch equation:

$$pH = pK_a + \log\left(\frac{[A^-]}{[HA]}\right)$$

Here, \(pK_a\) is derived from the relationship \(K_a = \frac{K_w}{K_b}\), where \(K_w\) is the ion product of water (1.0 × 10^-14) and \(K_b\) for ammonia is 1.76 × 10^-5. This calculation yields \(K_a\) as 5.68 × 10^-10. Plugging the values into the Henderson-Hasselbalch equation, the resulting pH is approximately 10.746.

As the titration progresses and more strong acid is added, the pH decreases from an initial value of just over 11, demonstrating the expected trend of pH reduction as the equivalence point is approached. Understanding these calculations and the behavior of the system before and at the equivalence point is essential for mastering acid-base titrations.

Weak Base-Strong Acid Titrations

In a titration involving a weak base and a strong acid, the equivalence point is characterized by the complete neutralization of the reactants, resulting in an acidic solution. Specifically, when titrating ammonia (a weak base) with hydrochloric acid (a strong acid), the pH at the equivalence point will be less than 7 at 25 degrees Celsius due to the formation of the conjugate acid, ammonium ion.

For example, consider a titration where 100 mL of 0.100 M ammonia is titrated with 50 mL of 0.200 M HCl. To find the moles of each reactant, the volumes are converted to liters and multiplied by their molarities. This calculation reveals that both reactants provide 0.100 moles, which completely neutralize each other, leaving 0 moles of ammonia and HCl, but producing 0.100 moles of ammonium ion.

To determine the pH at the equivalence point, it is essential to calculate the molarity of the resulting weak acid (ammonium ion). The total volume after mixing is 0.100 L + 0.050 L = 0.150 L. The molarity of the ammonium ion is calculated as:

$$\text{Molarity} = \frac{\text{moles of ammonium ion}}{\text{total volume}} = \frac{0.100 \text{ moles}}{0.150 \text{ L}} = 0.667 \text{ M}$$

Next, an ICE (Initial, Change, Equilibrium) chart is utilized to find the concentration of hydronium ions (H₃O⁺) produced from the dissociation of the weak acid. The dissociation constant (Kₐ) for ammonium ion is needed, which is given as 5.68 x 10^-10. The equilibrium expression can be set up as follows:

$$K_a = \frac{x^2}{[0.667 - x]}$$

Assuming x is small compared to 0.667, the equation simplifies to:

$$K_a \approx \frac{x^2}{0.667}$$

Solving for x gives:

$$x^2 = K_a \times 0.667 = 5.68 \times 10^{-10} \times 0.667$$

$$x^2 = 3.789 \times 10^{-10}$$

Taking the square root yields:

$$x = 1.95 \times 10^{-5}$$

This value represents the concentration of H₃O⁺ ions. To find the pH, the negative logarithm of the hydronium ion concentration is calculated:

$$\text{pH} = -\log(1.95 \times 10^{-5}) \approx 4.71$$

Thus, at the equivalence point of this titration, the pH is 4.71, confirming that the solution is indeed acidic due to the presence of the conjugate acid formed from the weak base and strong acid neutralization. As more strong acid is added, the pH will continue to decrease, illustrating the behavior of weak bases in strong acid titrations.

Weak Base-Strong Acid Titrations

In a titration involving a weak base and a strong acid, understanding the behavior after the equivalence point is crucial. After reaching the equivalence point, where the amount of strong acid, such as hydrochloric acid (HCl), equals the amount of weak base, any additional strong acid will remain unreacted. For instance, if 50 mL of HCl is required to reach the equivalence point and we then add an additional 10 mL, we have a total of 60 mL of HCl.

To analyze the situation, we can use an ICF (Initial, Change, Final) table. First, convert the volumes from milliliters to liters by dividing by 1,000. Then, multiply the volume in liters by the molarity of the solutions to find the number of moles. In this case, the moles of the strong acid will exceed those of the weak base, leading to a calculation where the moles of the weak base are subtracted from the moles of the strong acid. This results in no remaining weak base and a surplus of strong acid.

To find the concentration of the remaining strong acid, we use the formula:

C = \frac{n}{V}

where C is the concentration, n is the number of moles of HCl left, and V is the total volume in liters. In this example, if the moles of HCl remaining are calculated and the total volume is 0.100 L + 0.060 L = 0.160 L, the concentration of HCl can be determined.

Once the concentration is known, the pH can be calculated using the formula:

pH = -\log[H^+]

For a concentration of 0.125 M HCl, the pH would be:

pH = -\log(0.125) \approx 0.903

This illustrates the significant drop in pH from the initial value above 11 when only the weak base was present, to a pH just below 1 after the addition of excess strong acid. Understanding these steps and the use of ICF or ICE charts is essential for accurately determining the pH during different stages of a titration.

Weak Base-Strong Acid Titrations Calculations

In the titration of a weak base, such as methylamine (CH₃NH₂), with a strong acid like hydrochloric acid (HCl), it is essential to understand the underlying chemistry and calculations involved. Methylamine has a base dissociation constant (K_b) of 4.4 × 10⁻⁴, indicating that it is a weak base. When titrated with HCl, the reaction can be represented as:

CH₃NH₂ + HCl → CH₃NH₃⁺ + Cl⁻

In this scenario, we start with 50 mL of 0.150 M methylamine and 75 mL of 0.200 M hydrochloric acid. To find the pH after the titration, we first need to calculate the moles of each reactant. Converting the volumes from milliliters to liters gives:

For methylamine:
Moles = Volume (L) × Molarity (mol/L) = 0.050 L × 0.150 mol/L = 0.0075 moles

For hydrochloric acid:
Moles = Volume (L) × Molarity (mol/L) = 0.075 L × 0.200 mol/L = 0.015 moles

Since we have more moles of HCl than methylamine, the weak base will be completely consumed, leaving us with excess strong acid. The remaining moles of HCl can be calculated as:

Remaining moles of HCl = 0.015 moles (initial) - 0.0075 moles (reacted) = 0.0075 moles

Next, we determine the total volume of the solution after mixing, which is:

Total Volume = 50 mL + 75 mL = 125 mL = 0.125 L

Now, we can find the concentration of the remaining HCl:

Concentration of HCl = Remaining moles / Total volume = 0.0075 moles / 0.125 L = 0.060 M

To calculate the pH, we use the formula:

pH = -log[H⁺] = -log(0.060) ≈ 1.22

This pH value indicates that we are in the region after the equivalence point of the titration, where excess strong acid significantly influences the pH. Understanding these calculations and concepts is crucial for analyzing weak base-strong acid titrations effectively.

Weak Base-Strong Acid Titrations Calculations

In this scenario, we are tasked with calculating the pH of a solution formed by mixing sodium acetate and acetic acid with perchloric acid. The key components involved are 75 mL of 0.100 M sodium acetate, 75 mL of 0.150 M acetic acid, and 0.0040 moles of perchloric acid. The dissociation constant (K_a) for acetic acid is given as 1.8 × 10⁻⁵.

When mixing these solutions, we recognize that perchloric acid, a strong acid, will react with the acetate ion, which is the conjugate base of acetic acid. The reaction can be represented as:

CH₃COO⁻ + HClO₄ → CH₃COOH + ClO₄⁻

To analyze the reaction, we can use an ICF (Initial, Change, Final) chart, focusing on the moles of each species involved. First, we convert the volumes of the solutions to moles:

• Sodium acetate: 0.075 L × 0.100 mol/L = 0.0075 moles
• Acetic acid: 0.075 L × 0.150 mol/L = 0.01125 moles
• Perchloric acid: 0.0040 moles (given)

Next, we analyze the reactants. The perchloric acid will react with the sodium acetate, consuming 0.0040 moles of acetate:

• Remaining acetate: 0.0075 moles - 0.0040 moles = 0.0035 moles
• Acetic acid will increase by the same amount of moles of perchloric acid consumed: 0.01125 moles + 0.0040 moles = 0.01525 moles

At the end of the reaction, we have a buffer solution consisting of the remaining conjugate base (acetate) and the weak acid (acetic acid). To find the pH of this buffer solution, we apply the Henderson-Hasselbalch equation:

pH = pK_a + log₁₀([A⁻]/[HA])

First, we calculate pK_a:

pK_a = -log₁₀(K_a) = -log₁₀(1.8 × 10⁻⁵) ≈ 4.74

Now, substituting the values into the Henderson-Hasselbalch equation:

pH = 4.74 + log₁₀(0.0035/0.01525)

Calculating the log term gives us:

log₁₀(0.0035/0.01525) ≈ -0.43

Thus, the final pH is:

pH ≈ 4.74 - 0.43 = 4.31

In conclusion, the pH of the resulting solution after mixing these components is approximately 4.31, indicating that we are still before the equivalence point in this titration, as a buffer is present.

Here’s what students ask on this topic:

What is the equivalence point in a weak base-strong acid titration?

The equivalence point in a weak base-strong acid titration occurs when the moles of the strong acid (titrant) added equal the moles of the weak base (analyte) initially present. At this point, the weak base has been completely neutralized by the strong acid, resulting in the formation of a conjugate acid. The pH at the equivalence point is less than 7 because the solution is acidic. This is due to the presence of the conjugate acid, which partially dissociates in water to produce H₃O⁺ ions, lowering the pH.

How do you calculate the pH before the equivalence point in a weak base-strong acid titration?

Before the equivalence point, the pH is calculated using an ICF (Initial, Change, Final) chart to determine the moles of the weak base and the strong acid. The remaining weak base and its conjugate acid form a buffer solution. The pH can be calculated using the Henderson-Hasselbalch equation:

$\mathrm{pH}=pKa+\frac{\log }{\frac{\mathrm{[A^-]}}{\mathrm{[HA]}}}$

where [A^-] is the concentration of the conjugate base and [HA] is the concentration of the weak acid. This equation helps determine the pH of the buffer solution formed during the titration.

Why is the pH at the equivalence point of a weak base-strong acid titration less than 7?

The pH at the equivalence point of a weak base-strong acid titration is less than 7 because the solution contains the conjugate acid of the weak base. When the weak base is completely neutralized by the strong acid, the resulting solution contains the conjugate acid, which partially dissociates in water to produce H₃O⁺ ions. This increase in H₃O⁺ concentration lowers the pH, making the solution acidic. The exact pH can be calculated using an ICE chart and the acid dissociation constant (K_a) of the conjugate acid.

How do you calculate the pH after the equivalence point in a weak base-strong acid titration?

After the equivalence point, the pH is determined by the excess strong acid remaining in the solution. First, calculate the moles of excess HCl by subtracting the moles of the weak base from the total moles of HCl added. Then, find the concentration of the excess HCl by dividing the moles of excess HCl by the total volume of the solution. Finally, calculate the pH using the formula:

$\mathrm{pH}=-\log \left(\mathrm{[H⁺]}\right)$

where [H⁺] is the concentration of the excess HCl. This gives the pH of the solution after the equivalence point.

What is the role of the Henderson-Hasselbalch equation in weak base-strong acid titrations?

The Henderson-Hasselbalch equation is used to calculate the pH of buffer solutions formed during a weak base-strong acid titration. When a weak base and its conjugate acid are present in significant amounts, they form a buffer solution. The equation is:

$\mathrm{pH}=pKa+\frac{\log }{\frac{\mathrm{[A^-]}}{\mathrm{[HA]}}}$

where [A^-] is the concentration of the conjugate base and [HA] is the concentration of the weak acid. This equation helps determine the pH of the buffer solution, which is crucial for understanding the pH changes during the titration process.