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If 0.20 mol/L CH3COOH and 0.50 mol/L CH3COO– together make a buffer solution, calculate the pH of the solution if the acid dissociation constant of CH3COOH is 1.8 × 10-5.
- a)2.09
- b)5.14
- c)2.65
- d)3.98
Correct answer is option 'B'. Can you explain this answer?
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Most Upvoted Answer
If 0.20 mol/L CH3COOH and 0.50 mol/L CH3COO– together make a buffer so...
Understanding Buffer Solutions
Buffer solutions resist changes in pH when small amounts of acid or base are added. In this case, we have acetic acid (CH3COOH) and its conjugate base, acetate (CH3COO–), forming a buffer.
Given Data
- Concentration of CH3COOH = 0.20 mol/L
- Concentration of CH3COO– = 0.50 mol/L
- Acid dissociation constant (Ka) of CH3COOH = 1.8 × 10⁻⁵
Calculating pKa
To find pH, we first need to calculate the pKa from the given Ka:
- pKa = -log(Ka)
- pKa = -log(1.8 × 10⁻⁵) ≈ 4.74
Using the Henderson-Hasselbalch Equation
The pH of a buffer solution can be calculated using the Henderson-Hasselbalch equation:
Where:
- [A⁻] = concentration of the conjugate base (CH3COO–)
- [HA] = concentration of the weak acid (CH3COOH)
Substituting the Values
Now we substitute in the values:
- pH = 4.74 + log(0.50 / 0.20)
Calculating the Logarithm
- Log(0.50 / 0.20) = log(2.5) ≈ 0.39794
Final pH Calculation
Now, substituting this back into the equation:
- pH = 4.74 + 0.39794 ≈ 5.14
Conclusion
The calculated pH of the buffer solution is approximately 5.14, which corresponds to option 'B'.
Buffer solutions resist changes in pH when small amounts of acid or base are added. In this case, we have acetic acid (CH3COOH) and its conjugate base, acetate (CH3COO–), forming a buffer.
Given Data
- Concentration of CH3COOH = 0.20 mol/L
- Concentration of CH3COO– = 0.50 mol/L
- Acid dissociation constant (Ka) of CH3COOH = 1.8 × 10⁻⁵
Calculating pKa
To find pH, we first need to calculate the pKa from the given Ka:
- pKa = -log(Ka)
- pKa = -log(1.8 × 10⁻⁵) ≈ 4.74
Using the Henderson-Hasselbalch Equation
The pH of a buffer solution can be calculated using the Henderson-Hasselbalch equation:
Where:
- [A⁻] = concentration of the conjugate base (CH3COO–)
- [HA] = concentration of the weak acid (CH3COOH)
Substituting the Values
Now we substitute in the values:
- pH = 4.74 + log(0.50 / 0.20)
Calculating the Logarithm
- Log(0.50 / 0.20) = log(2.5) ≈ 0.39794
Final pH Calculation
Now, substituting this back into the equation:
- pH = 4.74 + 0.39794 ≈ 5.14
Conclusion
The calculated pH of the buffer solution is approximately 5.14, which corresponds to option 'B'.
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Community Answer
If 0.20 mol/L CH3COOH and 0.50 mol/L CH3COO– together make a buffer so...
We have the Henderson-Hasselbalch equation as pH = pKa + log[salt]/[acid]. So by substituting the concentrations of salt and acid along with the acid dissociation constant, we get pH = -log[1.8 × 10-5] + log [0.50 mol/L]/[0.20 mol/L] = 5.14.
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If 0.20 mol/L CH3COOH and 0.50 mol/L CH3COO– together make a buffer solution, calculate the pH of the solution if the acid dissociation constant of CH3COOH is 1.8 × 10-5.a)2.09b)5.14c)2.65d)3.98Correct answer is option 'B'. Can you explain this answer?
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