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Which of the following statements are incorrect?

a)  pKw = 14

b)  pKw = pKa + pKb

c)  Kw = 1 x 10-14

d)  Kw = Ka + Kb


	Only d is incorrect
	Only b is incorrect
	All are correct
	Both a and b are incorrect
	Both c and d are incorrect

The statements can be evaluated as follows:

a) **pKw = 14**  
This statement is correct at 25°C, where the ion product of water, Kw, equals \(1 \times 10^{-14}\).

b) **pKw = pKa + pKb**  
This statement is correct. It reflects the relationship between the pKw, pKa of the acid, and pKb of the conjugate base.

c) **Kw = \(1 \times 10^{-14}\)**  
This statement is correct at 25°C.

d) **Kw = Ka + Kb**  
This statement is incorrect. The relationship between Ka, Kb, and Kw is not additive; rather, it is multiplicative: \(K_w = K_a \cdot K_b\).

Based on this evaluation, the correct choice is:

**Only d is incorrect.**

The background material for this experiment defines Ka as:

(select all that apply)



	
The Henderson-Hasselbalch constant

	
The equilibrium constant for a weak acid in water

	
The acid ionization constant

The correct definitions for \( K_a \) are:

- **The equilibrium constant for a weak acid in water**
- **The acid ionization constant**

So, the correct selections from the options provided would be:

- **The equilibrium constant for a weak acid in water**
- **The acid ionization constant**

**Note:** The Henderson-Hasselbalch equation relates the pH of a buffer solution to the concentrations of an acid and its conjugate base, but it does not define \( K_a \) itself. Therefore, **"The Henderson-Hasselbalch constant"** is not correct in the context of defining \( K_a \).

By definition, pOH = - log [OH-]


	True
	False

**True.**

By definition, \( \text{pOH} = -\log [\text{OH}^-] \). This equation describes the relationship between the hydroxide ion concentration \([\text{OH}^-]\) and the pOH value, similar to how pH relates to the hydrogen ion concentration.

In a titration of a strong base with a weak acid at 25oC, the theoretical pH at equivalence point is: 

	
 less than 7.00

	
 exactly 7.00

	
 greater than 7.00

	
 not enough information is given to determine

In a titration of a strong base with a weak acid, the theoretical pH at the equivalence point is:

**greater than 7.00.**

At the equivalence point, all the weak acid has reacted with the strong base to form its conjugate base, which will hydrolyze in water to produce hydroxide ions (OH⁻). This results in a pH that is above 7, typically greater than 7, due to the basic nature of the conjugate base formed from the weak acid.

Which of the following statements are true regarding an acid-base indicator?

(Select all that apply)


	
  its colour changes at the equivalence point of any acid-base titration

	
 its colour change indicates pH of the solution has reached 7.00 

	
  it is itself a weak acid or a weak base

	
  its colour changes when it gains or loses a H+ ion

	
  its various protonated species have different colours

The true statements regarding an acid-base indicator are:

1. **It is itself a weak acid or a weak base.**  
   Indicators typically are weak acids or weak bases that can undergo ionization.

2. **Its colour changes when it gains or loses a H⁺ ion.**  
   Acid-base indicators change colour based on their protonation state, which involves gaining or losing a hydrogen ion.

3. **Its various protonated species have different colours.**  
   Many indicators exist in different forms depending on their protonation state, and these forms often have different colours.

The other statements are not generally true:

- **Its colour changes at the equivalence point of any acid-base titration.** (False - The colour change occurs at a certain pH range; it may not coincide with the equivalence point in all cases.)

- **Its colour change indicates the pH of the solution has reached 7.00.** (False - Indicators change colour at specific pH ranges, which may be above or below 7.00 depending on the indicator used.)

So, the selections that are true are:

- **It is itself a weak acid or a weak base.**
- **Its colour changes when it gains or loses a H⁺ ion.**
- **Its various protonated species have different colours.**

What is the pH transition range of methyl red and what colour change will be observed over these values? 

	
  6.0 – 8.0, red → blue

	
  4.8 – 6.0, red → yellow

	
  8.0 - 9.6, colourless → pink

	
  8.0 – 9.6, pink → colourless 

	
  4.8 – 6.0, yellow → red 

The correct pH transition range of methyl red is:

**4.8 – 6.0, red → yellow.**

In this pH range, methyl red changes from red in acidic conditions to yellow in more neutral to slightly alkaline conditions.

What is the pH transition range of phenolphthalein and what colour change will be observed over these values? 

	
  6.0 – 8.0, red → blue

	
  4.8 – 6.0, yellow → red 

	
  8.0 – 9.6, pink → colourless 

	
  8.0 - 9.6, colourless → pink

	
  4.8 – 6.0, red → yellow

The correct pH transition range of phenolphthalein is:

**8.0 - 9.6, colourless → pink.**

In this pH range, phenolphthalein changes from colourless in acidic and neutral solutions to pink in basic solutions.

Acetic acid is a weak acid with the formula CH3COOH; the Ka for acetic acid is 1.76 x 10-5

In aqueous solution, acetic acid partially dissociates according to the following reaction:

CH3COOH + H2O ⇔ CH3COO- + H3O+

Calculate the pH of the acetic acid solution described below:

Volume:  375 mL

Concentration:  0.1448 M

Since this is a weak acid, you can assume the amount of acid dissociated is << 5% of the total amount of acid present.

 

	
2.57

	
1.27

	
2.80

	
3.01

	
0.84

To calculate the pH of a 0.1448 M solution of acetic acid (\( \text{CH}_3\text{COOH} \)), we can set up the equilibrium expression based on its dissociation:

\[
\text{CH}_3\text{COOH} + \text{H}_2\text{O} \rightleftharpoons \text{CH}_3\text{COO}^- + \text{H}_3\text{O}^+
\]

The equilibrium constant for this reaction is given by:

\[
K_a = \frac{[\text{CH}_3\text{COO}^-][\text{H}_3\text{O}^+]}{[\text{CH}_3\text{COOH}]}
\]

1. **Define variables**:
    - Initial concentration of \( \text{CH}_3\text{COOH} \): \( C = 0.1448 \, \text{M} \)
    - Let \( x \) be the amount of acetic acid that dissociates.

At equilibrium:
- \([\text{CH}_3\text{COOH}] = C - x \approx 0.1448 \, \text{M}\) (since \( x \) will be small)
- \([\text{CH}_3\text{COO}^-] = x\)
- \([\text{H}_3\text{O}^+] = x\)

2. **Set up the equilibrium expression**:
\[
K_a = \frac{x \cdot x}{C - x} = \frac{x^2}{0.1448 - x}
\]

Since \( K_a \) is small and \( x \) is assumed to be much smaller than 0.1448, we will simplify this to:
\[
K_a \approx \frac{x^2}{0.1448}
\]

3. **Plug in the values**:
\[
1.76 \times 10^{-5} = \frac{x^2}{0.1448}
\]

4. **Solve for \( x \)**:
\[
x^2 = 1.76 \times 10^{-5} \times 0.1448
\]
\[
x^2 = 2.54848 \times 10^{-6}
\]
\[
x = \sqrt{2.54848 \times 10^{-6}} \approx 0.001597 \, \text{M}
\]

5. **Calculate pH**:
\[
[\text{H}_3\text{O}^+] \approx x \approx 0.001597 \, \text{M}
\]
\[
\text{pH} = -\log(0.001597) \approx 2.80
\]

Therefore, the calculated pH of the acetic acid solution is:

**2.80**.

Pyridine is a weak base with the formula, C5H5N, the Kb for Pyridine is 1.8 x 10-9

In aqueous solution, pyridine partially dissociates according to the following reaction:

C5H5N + H2O ⇔ C5H5NH+ + OH-

Use the Kb equation to calculate the pH of the aqueous pyridine solution described below:

Volume:  400 mL

Concentration:  0.2851 M

Since this is a weak base, you can assume the amount of base dissociated is << 5% of the total amount of base present.

Report your answer with the correct number of significant digits.

Number
 

 

To calculate the pH of a 0.2851 M solution of pyridine (\( \text{C}_5\text{H}_5\text{N} \)), we can set up the equilibrium expression based on its dissociation:

\[
\text{C}_5\text{H}_5\text{N} + \text{H}_2\text{O} \rightleftharpoons \text{C}_5\text{H}_5\text{NH}^+ + \text{OH}^-
\]

The base dissociation constant (\( K_b \)) for this reaction is given by:

\[
K_b = \frac{[\text{C}_5\text{H}_5\text{NH}^+][\text{OH}^-]}{[\text{C}_5\text{H}_5\text{N}]}
\]

1. **Define Variables**:
   - Initial concentration of \( \text{C}_5\text{H}_5\text{N} \): \( C = 0.2851 \, \text{M} \)
   - Let \( x \) be the amount of pyridine that dissociates.

At equilibrium:
- \([\text{C}_5\text{H}_5\text{N}] = C - x \approx 0.2851 \, \text{M}\) (since \( x \) will be small)
- \([\text{C}_5\text{H}_5\text{NH}^+] = x\)
- \([\text{OH}^-] = x\)

2. **Set up the \( K_b \) expression**:
\[
K_b = \frac{x \cdot x}{C - x} = \frac{x^2}{0.2851 - x}
\]

Since \( K_b \) is small and \( x \) is assumed to be much smaller than 0.2851, we can simplify this to:
\[
K_b \approx \frac{x^2}{0.2851}
\]

3. **Plug in the values**:
\[
1.8 \times 10^{-9} = \frac{x^2}{0.2851}
\]

4. **Solve for \( x \)**:
\[
x^2 = 1.8 \times 10^{-9} \times 0.2851
\]
\[
x^2 = 5.13018 \times 10^{-10}
\]
\[
x = \sqrt{5.13018 \times 10^{-10}} \approx 2.27 \times 10^{-5} \, \text{M}
\]

5. **Calculate the \( [\text{OH}^-] \)**:
Since \( [\text{OH}^-] \) is approximately equal to \( x \):
\[
[\text{OH}^-] = 2.27 \times 10^{-5} \, \text{M}
\]

6. **Calculate the pOH**:
\[
\text{pOH} = -\log(2.27 \times 10^{-5}) \approx 4.44
\]

7. **Calculate the pH**:
Using the relationship \( \text{pH} + \text{pOH} = 14 \):
\[
\text{pH} = 14 - \text{pOH} = 14 - 4.44 \approx 9.56
\]

Therefore, the pH of the 0.2851 M pyridine solution is approximately **9.56** when reported to four significant figures.

In an acid base titration, the equivalence point is the point at which:

	
 the pH value starts to change rapidly

	
 the pH value has reached a plateau

	
 all of the analyte present has reacted with the amount of titrant added

	
 the indicator changes colour

	
 the pH value has reached 7.00

In an acid-base titration, the equivalence point is the point at which:

**all of the analyte present has reacted with the amount of titrant added.**

At the equivalence point, the number of moles of acid is equal to the number of moles of base in the reaction, meaning that the reaction is complete. While it is true that the pH may change rapidly around this point and that an indicator may change color, the defining characteristic of the equivalence point is the complete reaction of the analyte with the titrant. The pH at the equivalence point does not have to be exactly 7.00, especially in the case of weak acid-strong base or weak base-strong acid titrations, where the pH will be greater or less than 7, respectively.