A chemistry graduate student is given 250. mL of a 1.40M ammonia (NH3) solution. Ammonia is a weak base with K-1.8x10should the student dissolve in the NH3 solution to turn it into a buffer with pH = 9.17?What mass of NH, BrYou may assume that the volume of the solution doesn't change when the NH4Br is dissolved in it. Be sure your answer has a unit symbol, and round it to 2significant digits.?

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Answered: A chemistry graduate student is given…

Chemistry

10th Edition

ISBN:9781305957404

Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Chapter1: Chemical Foundations

Section: Chapter Questions

Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...

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A) Determine the pH of the solution that results from the mixing of 60.0 mL of 0.100 M NaOH, 94.0 mL of 0.0500 M KOH, 62.5 mL of 0.075 M HCl, 37.0 mL of 0.065 M HNO3, and 3.00 quarts of distilled water. B) Calculate the pH of a titration of 50.00 mL of 0.100 M Phenylacetic acid, Ka = 4.9 x 10-5, with 0.100 M NaOH at the following points: a. Before any NaOH is added. b. After 18.7 mL of NaOH are added. c. After 25.00 mL of NaOH are added. d. After 50.00 mL of NaOH are added. e. After 53.00 mL of NaOH are added. C) What is the pKa of the acid (show calculation)?
Suppose a student performed a similar standardization titration experiment. Below is the calculated concentration of NaOH from each of their titrations. Titration #1 Titration #2 Titration #3 Concentration of NaOH (M) 0.9554 0.9540 0.9551 Using the 3 concentrations, calculate the average concentration of the student's NaOH solution and calculate the ± error in parts per thousand (ppt). Show the complete calculation. Express your final answer in the form of average (with units) ± error in ppt. use the right significant figures
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In an advanced chemistry lab, a student prepared solutions of 0.1 M citric acid (H3C6H5O7), 0.1 M lactic acid (HC3H5O3) and 0.1 M sulfuric acid (H2SO4). Unfortunately, this advanced student neglected to label the samples of these acids! But all is not lost! The student recognizes that one of these is a monoprotic, one is a diprotic, and the other a triprotic acid. To determine the identity of the solutions, he decided to titrate the acids with 0.0500 M NAOH. Use the data he collected for his first titration to determine the identity of that acid. Volume of Acid Titrated: 10.00 mL Molarity NaOH: 0.0500 M Final Buret Reading: 43.12 mL Initial Buret Reading: 2.98 m Volume of NaOH used This acid is
A solution is made by dissolving 15.1 g of CH₃NH₃NO₃ in 500.0 mL of water. Using Ka(CH₃NH₃⁺) = 2.3 × 10⁻¹¹, determine the pH of the solution.
An analytical chemist weighs out 0.109 g of an unknown triprotic acid into a 250 mL volumetric flask and dilutes to the mark with distilled water. She then titrates this solution with 0.0800M NAOH solution. When the titration reaches the equivalence point, the chemist finds she has added 41.7 mL of NaOH solution. Calculate the molar mass of the unknown acid. Be sure your answer has the correct number of significant digits. x10 mol
In the biochemistry lab, a fellow student left four solutions out on the bench before rushing to class. The student had labeled each solution with its hydrogen ion concentration, [H*]. The professor asks you to put the solutions away in the acid-base cabinet. You remember from general chemistry that acids and bases must be stored separately to prevent corrosive reactions from occurring within the cabinet. Therefore, you first need to determine the pH of each solution. Calculate the pH of solution #1, which has an [H+] of 0.47 M. Calculate the pH of solution #2, which has an [H+] of 3.6 × 10-13 M. Calculate the pH of solution #3, which has an [H+] of 2.1 x 10-10 M. Calculate the pH of solution #4, which has an [H*] of 0.00060 M. solution #1 pH: solution #2 pH: solution #3 pH: solution #4 pH:
complete this precipitation reaction step by step _______K2CO3 (aq) + _______ Ca(Cl2H5O2)2 (aq) -------->
The arsenate sample was treated using silver ion to form a precipitate, silver arsenate. It took exactly 25.0 mL of a silver ion solution to remove all of the arsenate. The concentration of the silver ion solution is given below. Calculate the number of moles of silver ion used. 0.102 M Ag+. B) .Calculate the number of moles of arsenic in the precipitate. The formula for the precipitate is given below. Ag3AsO4
Help with the following question ( use kg for unit)
Consider the following reactions: Reaction 1: HCIO, + KOH → KC1O, + H;O Reaction 2: HC;H;O2 + KOH → KC;H;O2+ H;O In both reactions, 15.00 mL of 0.8116 M acid is titrated with 0.4217 M KOH. A volume of 28.87 mL of KOH is required to reach the equivalence point in both reactions. a) What is the pH of the HC1O, solution before the titration has started? (No KOH has been added.) b) What is the initial pH of HC;H;O2 solution before the titration has started? K= 6.46 x 10-5 (Ok to use x is small.) c) What is the pH of Reaction 1 at the equivalence point? d) What is the pH of Reaction 2 at the equivalence point? e) What is the pH of Reaction 1 after 32.00 mL of KOH has been added? f) What is the pH of Reaction 2 after 18.00 mL of KOH has been added?
The pH scale was designed to make it convenient to express hydrogen ion concentrations that are small in aqueous solutions. The definiton of pH is in terms of base 10 logarithms. pH = - log[H*] where [H*] is the hydrogen ion concentration. a. If the hydrogen ion concentration in a solution is 7.26 x 10 mol L, the pH is b. If the pH of a solution is 3.607, the hydrogen ion concentration is mol L.

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