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Ch.17 - Aqueous Ionic Equilibrium
Chapter 17, Problem 116
A buffer is created by combining 3.55 g of NH3 with 4.78 g of HCl and diluting to a total volume of 750.0 mL. Determine the pH of the buffer.
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Calculate the moles of NH3 using its molar mass: \( \text{moles of NH}_3 = \frac{3.55 \text{ g}}{17.03 \text{ g/mol}} \).
Calculate the moles of HCl using its molar mass: \( \text{moles of HCl} = \frac{4.78 \text{ g}}{36.46 \text{ g/mol}} \).
Determine the limiting reactant by comparing the moles of NH3 and HCl, and calculate the moles of NH4+ formed.
Use the Henderson-Hasselbalch equation: \( \text{pH} = \text{pK}_a + \log \left( \frac{[\text{base}]}{[\text{acid}]} \right) \), where \( \text{pK}_a \) is for NH4+ and the concentrations are based on the remaining NH3 and formed NH4+.
Calculate the concentrations of NH3 and NH4+ by dividing their moles by the total volume of the solution (0.750 L) and substitute these values into the Henderson-Hasselbalch equation to find the pH.
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Textbook Question
A 1.0-L buffer solution initially contains 0.25 mol of NH3 and 0.25 mol of NH4Cl. In order to adjust the buffer pH to 8.75, should you add NaOH or HCl to the buffer mixture? What mass of the correct reagent should you add?
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To adjust the pH of a 250.0-mL buffer solution initially containing 0.025 mol of HCHO2 and 0.025 mol of NaCHO2 to 4.10, should you add NaOH or HCl, and what mass of the correct reagent should you add?
Textbook Question
In analytical chemistry, bases used for titrations must often be standardized; that is, their concentration must be precisely determined. Standardization of sodium hydroxide solutions can be accomplished by titrating potassium hydrogen phthalate (KHC8H4O4), also known as KHP, with the NaOH solution to be standardized. b. The titration of 0.5527 g of KHP required 25.87 mL of an NaOH solution to reach the equivalence point. What is the concentration of the NaOH solution?
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