Write the Henderson-Hasselbalch equation for a propanoic acid solution (CH;CH,CO,H, pKa=4.874) using the symbols HA and A, and the given pka value for propanoic acid in the expression. pH = + log - Using the equation to calculate the quotient [A-]/[HA] at three different pH values. pH = 5.321 ГА УНА] - pH = 4.079 [А УНА] 3 pH = 4.874 [А УНА] %3

Write the Henderson-Hasselbalch equation for a propanoic acid solution (CH;CH,CO,H, pKa=4.874) using the symbols HA and A, and the given pka value for propanoic acid in the expression. pH = + log - Using the equation to calculate the quotient [A-]/[HA] at three different pH values. pH = 5.321 ГА УНА] - pH = 4.079 [А УНА] 3 pH = 4.874 [А УНА] %3

Chemistry: Principles and Practice

3rd Edition

ISBN:9780534420123

Author:Daniel L. Reger, Scott R. Goode, David W. Ball, Edward Mercer

Publisher:Daniel L. Reger, Scott R. Goode, David W. Ball, Edward Mercer

Chapter16: Reactions Between Acids And Bases

Section: Chapter Questions

Problem 16.81QE: Write the chemical equation and the expression for the equilibrium constant, and calculate Kb for...

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The pH of a 0.0200 M solution of an unknown acid is 2.56. What is the Ka of this acid? To solve this problem: Write the acid dissociation equilibrium for the generic acid “HA” Set up an ICE chart ( with x = the concentration of H3O+ at equilibrium) Write the expression for Ka. Fill this in with the concentrations of H3O+, A- and HA at equilibrium, in terms of x. What is x? Can you find it from the given information? You should be able to use the pH to get the concentration of H3O+, which is x. Do this. Now that you know x, plug in into the Ka expression and find Ka.
Determine the pH of a solution by constructing a BCA table, constructing an ICE table, writing the equilibrium constant expression, and using this information to determine the pH. Complete Parts 1-4 before submitting your answer. NEXT > 0.040 mol of solid NaOH is added to a 200.0 mL buffer containing 0.100 mol C6H5NH3Cl and 0.500 M C6H5NH₂. Fill in the table with the appropriate value for each involved species to determine the moles of reactant and product after the reaction of the acid and base. You can ignore the amount of water in the reaction. Before (mol) Change (mol) After (mol) 0.060 Initial (M) Change (M) Equilibrium (M) 0 -X 0.700 - x 1 -0.060 C6H5NH3*(aq) + OH (aq) 0 1 0.040 0.140 0.500 + x CoHşNHz(aq) + 0.300 + x 2 0.500 0.500 - x -0.040 0.300 - x -0.140 Determine the pH of a solution by constructing a BCA table, constructing an ICE table, writing the equilibrium constant expression, and using this information to determine the pH. Complete Parts 1-4 before submitting your…
Calculate the change in pH that occurs when 1.10 mmol of a strong acid is added to 100. mL of the solutions listed below. Ka (CH3COOH) = = 1.75 x 10-5 a. 0.0680 M CH3COOH. Change in pH = b. 0.0680 M CH3COONa. Change in pH =
An organic acid (HA) has a molecular weight of 100g/mol, a Kow = 5.6 and a K, = 2.7 x 10-2. If originally 2.0 g of the acid is dissolved in 100 mL of octanol (there is no dissociation in octanol), which is then placed in contact with 100 mL of water, what will be the pH of the water? (Consider the equilibrium processes to be sequential and unrelated chemically) Possible answers: 1.89, 1.48, 2.01, 1.74, 1.61
The pH range of a solution can also be determined using indicators. Indicators are molecules that exhibit different colors over a specific pH range. Examples of indicators and the associated colors are shown below. pH Range and Colors 1 2 3 Indicator 4 6 7 Bromcresol green |Congo Red Methyl Yellow Thymol Blue Methyl Violet pH Range and Colors Indicator 2 3 4 5 7 Bromcresol green yellow Oblue Congo Red Methyl Yellow Thymol Blue Methyl Violet Jorange-red Dyellow Dyellow blue red red yellow Oviolet Bromcresol green, for example, would be green between pH 4-5, and gradually change to blue
Determine the pH of a solution by constructing a BCA table, constructing an ICE table, writing the equilibrium constant expression, and using this information to determine the pH. The value of Ka for HC₂H₂O₂ is 1.8 × 105. Complete Parts 1-4 before submitting your answer. 2 3 2 NEXT > 2 Two solutions are mixed: 20.0 mL of 0.20 M HC₂H₂O₂ and 20.0 mL of 0.10 M NaOH. Fill in the table with the appropriate value for each involved species to determine the moles of reactant and product after the reaction of the acid and base. You can ignore the amount of water in the reaction. Before (mol) Change (mol) After (mol) -2.0 × 10-³ 0 3.0 × 10-³ 1 HC,H,O,(aq) 20.0 -3.0 × 10-³ 2 0.20 OH(aq) 4.0 × 10-³ 0.10 3 -4.0 × 10-³ H₂O(1) 1.0 × 10-³ 4 -1.0 × 10³ C,H,O,(aq) RESET 2.0 × 10-³
A solution is prepared that is initially 0.45M in trimethylamine ((CH3)N), a weak base, and 0.37M in trimethylammonium bromide Complete the reaction table below, so that you could use it to calculate the pH of this solution. Use x to stand for the unknown change in [OH-]. You can leave out the M symbol for molarity. initial change final [(CH₂)₂N] [(CH₂)₂ 3 [(CH,),NH 0 0 0 0 0 [OH-] 0 0 0 ((CH3)₂NHBr
The Henderson-Hasselbalch equation relates the pH of a buffer solution to the pKa of its conjugate acid and the ratio of the concentrations of the conjugate base and acid. The equation is important in laboratory work that makes use of buffered solutions, in industrial processes where pH needs to be controlled, and in medicine, where understanding the Henderson-Hasselbalch equation is critical for the control of blood pH. Part A As a technician in a large pharmaceutical research firm, you need to produce 450. mL of 1.00 M potassium phosphate buffer solution of pH = 6.91. The pK₁ of H₂PO4¯ is 7.21. You have the following supplies: 2.00 L of 1.00 M KH₂PO4 stock solution, 1.50 L of 1.00 M K₂HPO4 stock solution, and a carboy of pure distilled H₂O. How much 1.00 M KH₂PO4 will you need to make this solution? Express your answer to three significant digits with the appropriate units. View Available Hint(s) 2 Volume of KH₂PO4 needed = Value Review | Constants | Periodic Table Submit μÃ Units ?
You are asked to prepare a pH = 3.00 buffer starting from 2.00 L of 0.025 M solution of benzoic acid (C6H5COOH). (a) What is the pH of the benzoic acid solution prior to adding sodium benzoate? (hint: write the reaction equation for the acid dissociation and then use the equilibrium constant expression to calculate [H + ]) (b) How many grams of sodium benzoate should be added to prepare the buffer? Neglect the small volume change that occurs when the sodium benzoate is added.(hint: use the equilibrium constant expression to calculate [C6H5COO− ] in the buffer)
(a) Calculate the pH in the solution formed by adding 10.0 mL of 0.050 M NaOH to 40.0 mL of 0.0250 M benzoic acid (C6H5COOH, Ka = 6.3 * 10-5). (b) Calculate the pH in the solution formed by adding 10.0 mL of 0.100 M HCl to20.0 mL of 0.100 M NH3.
You have 1.00 L of 0.10-M formic acid, HCOOH, whose Ka = 3.0 x 10-4. You want to bubble into the formic acid solution sufficient HCl gas to decrease the pH of the formic acid solution by 1.0 pH unit. Calculate the volume of HCl (liters) that must be used at STP to bring about the desired change in pH. Assume no volume change has occurred in the solution due to the addition of HCl gas.
Just as pH is the negative logarithm of [H3O+], pKa is the negative logarithm of Ka, - log Ka The Henderson-Hasselbalch equation is used to calculate the pH of buffer solutions: How many grams of dry NH4C1 need to be added to 2.00 L of a 0.600 mol L-' solution of ammonia, NH3, to prepare a buffer solution that has a pH of 8.80? Kh for ammonia is 1.8 × 10–5. pKa base] pH = pKa +log acid Express your answer to two significant figures and include the appropriate units. Notice that the pH of a buffer has a value close to the pKa of the acid, differing only by the logarithm of the concentration ratio [base]/[acid]. • View Available Hint(s) ? X•10n Value mass of NHẠCI = Units