Textbook Question
(b) What is the most significant difference between the sulfides precipitated in group 2 and those precipitated in group 3?
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Ch.17 - Additional Aspects of Aqueous Equilibria
Chapter 17, Problem 85
Furoic acid (HC5H3O3) has a Ka value of 6.76 × 10^-4 at 25 _x001F_C. Calculate the pH at 25 _x001F_C of (a) a solution formed by adding 25.0 g of furoic acid and 30.0 g of sodium furoate (NaC5H3O3) to enough water to form 0.250 L of solution; (b) a solution formed by mixing 30.0 mL of 0.250 M HC5H3O3 and 20.0 mL of 0.22 M NaC5H3O3 and diluting the total volume to 125 mL; (c) a solution prepared by adding 50.0 mL of 1.65 M NaOH solution to 0.500 L of 0.0850 M HC5H3O3.
Verified step by step guidance1
Step 1: For part (a), calculate the molarity of furoic acid (HC5H3O3) and sodium furoate (NaC5H3O3) in the solution. Use the formula: Molarity = (mass of solute / molar mass of solute) / volume of solution in liters.
Step 2: Recognize that the solution in part (a) is a buffer solution. Use the Henderson-Hasselbalch equation: pH = pKa + log([A-]/[HA]), where [A-] is the concentration of the conjugate base (sodium furoate) and [HA] is the concentration of the acid (furoic acid).
Step 3: For part (b), calculate the initial moles of HC5H3O3 and NaC5H3O3 using their given concentrations and volumes. Then, find their concentrations in the final 125 mL solution.
Step 4: Again, use the Henderson-Hasselbalch equation for part (b) to find the pH, using the concentrations calculated in Step 3.
Step 5: For part (c), calculate the moles of NaOH added and determine the moles of HC5H3O3 remaining after the reaction with NaOH. Use the stoichiometry of the reaction to find the concentrations of the remaining acid and its conjugate base, then apply the Henderson-Hasselbalch equation to find the pH.
Related Practice
Textbook Question
Derive an equation similar to the Henderson–Hasselbalch equation relating the pOH of a buffer to the pKb of its base component.
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Textbook Question
Rainwater is acidic because CO21g2 dissolves in the water, creating carbonic acid, H2CO3. If the rainwater is too acidic, it will react with limestone and seashells (which are principally made of calcium carbonate, CaCO3). Calculate the concentrations of carbonic acid, bicarbonate ion 1HCO3-2 and carbonate ion 1CO32 - 2 that are in a raindrop that has a pH of 5.60, assuming that the sum of all three species in the raindrop is 1.0 * 10-5 M.
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Textbook Question
The acid–base indicator bromcresol green is a weak acid. The yellow acid and blue base forms of the indicator are present in equal concentrations in a solution when the pH is 4.68. What is the pKa for bromcresol green?
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Open Question
Equal quantities of 0.010 M solutions of an acid HA and a base B are mixed. The pH of the resulting solution is 9.2. (a) Write the chemical equation and equilibrium-constant expression for the reaction between HA and B. (b) If Ka for HA is 8.0 × 10⁻⁵, what is the value of the equilibrium constant for the reaction between HA and B?
Textbook Question
Two buffers are prepared by adding an equal number of moles of formic acid (HCOOH) and sodium formate (HCOONa) to enough water to make 1.00 L of solution. Buffer A is prepared using 1.00 mol each of formic acid and sodium formate. Buffer B is prepared by using 0.010 mol of each. (b) Which buffer will have the greater buffer capacity?
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