Decomposition of Baking SodaPartial Pressure of CO₂The equilibrium constant, Kp = 0.35. The0.250 moles of NaHCO3 are heated at 145 °Cin an evacuated 2.00 L flask.+ H₂O(g)2NaHCO3(s) Na₂CO3(s) + CO₂(g) +=Calculate the partial pressure of CO2 atequilibrium.Copyright © 2003-2023 International Academy of Science. All Rights Reserved.What is the partial pressure ofCO2 at equilibrium?A. 0.592 atmB. 1.18 atmC. 4.29 atmD. 10.2 atm

Answered: Image /qna-images/answer/b2ce400c-1c28-41de-8f1e-4dcde96cc30e.jpg

Decomposition of Baking Soda Partial Pressure of CO₂ The equilibrium constant, Kp = 0.35. The 0.250 moles of NaHCO3 are heated at 145 °C in an evacuated 2.00 L flask. 2NaHCO3(s) ⇒ Na₂CO3(s) + -CO₂(g) + H₂O(g) Calculate the partial pressure of CO2 at equilibrium. What is the partial pressure of CO2 at equilibrium? A. 0.592 atm B. 1.18 atm C. 4.29 atm D. 10.2 atm

Decomposition of Baking Soda Partial Pressure of CO₂ The equilibrium constant, Kp = 0.35. The 0.250 moles of NaHCO3 are heated at 145 °C in an evacuated 2.00 L flask. 2NaHCO3(s) ⇒ Na₂CO3(s) + -CO₂(g) + H₂O(g) Calculate the partial pressure of CO2 at equilibrium. What is the partial pressure of CO2 at equilibrium? A. 0.592 atm B. 1.18 atm C. 4.29 atm D. 10.2 atm

Chemistry for Engineering Students

4th Edition

ISBN:9781337398909

Author:Lawrence S. Brown, Tom Holme

Publisher:Lawrence S. Brown, Tom Holme

Chapter12: Chemical Equilibrium

Section: Chapter Questions

Problem 12.76PAE: In a particular experiment, the equilibrium constant measured for the reaction,...

See similar textbooks

Similar questions

Write the expression for the equilibrium constant and calculate the partial pressure of CO2(g), given that Kp is 0.25 (at 427 C) for NaHCO3(s) NaOH(s) + CO2(g)
The following question is taken from a Chemistry Advanced Placement Examination and is used with the permission of the Educational Testing Service. Solve the following problem: MgF2(s)Mg2+(aq)+2F(aq) In a saturated solution of MgF2 at 18 C, the concentration of Mg2+ is 1.21103M . The equilibrium is represented by the preceding equation. (a) Write the expression for the solubility-product constant, Ksp, and calculate its value at 18 C. (b) Calculate the equilibrium concentration of Mg2+ in 1.000 L of saturated MgF2 solution at 18 C to which 0.100 mol of solid KF has been added. The KF dissolves completely. Assume the volume change is negligible. (c) Predict whether a precipitate of MgF2 will form when 100.0 mL of a 3.00103 -M solution of Mg(NO3)2 is mixed with 200.0 mL of a .2.00103 -M solution of NaF at 18 C. Show the calculations to support your prediction.. (d) At 27 C the concentration of Mg2+ in a saturated solution of MgF2 is 1.17103M . Is the dissolving of MgF2 in water all endothermic or an exothermic process? Give an explanation to support your conclusion.
. The solubility product of iron(III) hydroxide is very small: Ksp=41038at 25 °C. A classical method of analysis for unknown samples containing iron is to add NaOH or NH3. This precipitates Fe(OH)3, which can then be filtered and weighed. To demonstrate that the concentration of iron remaining in solution in such a sample is very small, calculate the solubility of Fe(OH)3in moles per liter and in grams per liter.
. K for copper(II)hydroxide, Cu(OH)2, has a value 2.21020at 25 °C. Calculate the solubility of copper(II) hydroxide in mol/L and g/L at 25 °C.
Use the solubility product constant from Appendix F to determine whether a precipitate will form if 10.0 mL of 1.0 106 M iron(II) chloride is added to 20.0 mL of 3.0 104 M barium hydroxide.
The Handbook of Chemistry and Physics (http://openstaxcollege.org/l/16Handbook) gives solubilities of the following compounds in grams per 100 mL of water. Because these compounds are only slightly soluble, assume that the volume does not change on dissolution and calculate the solubility product for each. (a) BaSiF6, 0.026 g/100 mL (contains SiF62- ions) (b) Ce(IO3)4, 1.5102 g/100 mL (c) Gd2(SO4)3, 3.98 g/100 mL (d) (NH4)2PtBr6, 0.59 g/100 mL (contains PtBr62- ions)
Because calcium carbonate is a sink for CO32- in a lake, the student in Exercise 12.39 decides to go a step further and examine the equilibrium between carbonate ion and CaCOj. The reaction is Ca2+(aq) + COj2_(aq) ** CaCO,(s) The equilibrium constant for this reaction is 2.1 X 10*. If the initial calcium ion concentration is 0.02 AI and the carbonate concentration is 0.03 AI, what are the equilibrium concentrations of the ions? A student is simulating the carbonic acid—hydrogen carbonate equilibrium in a lake: H2COj(aq) H+(aq) + HCO}‘(aq) K = 4.4 X 10"7 She starts with 0.1000 AI carbonic acid. What are the concentrations of all species at equilibrium?
What is the law of mass action? Is it true that the value of K depends on the amounts of reactants and products mixed together initially? Explain. Is it true that reactions with large equilibrium constant values are very fast? Explain. There is only one value of the equilibrium constant for a particular system at a particular temperature, but there is an infinite number of equilibrium positions. Explain.
In a particular experiment, the equilibrium constant measured for the reaction, Cl2(g)+NO2(g)Cl2NO2(g), is 2.8. Based on this measurement, calculate AG° for this reaction. Calculate AG° using data from Appendix E at the back of the book and discuss the agreement between your two calculations.
Because barium sulfate is opaque to X-rays, it is suspended in water and taken internally to make the gastrointestinal tract visible in an X-ray photograph. Although barium ion is quite toxic, barium sulfate’s /Csp of 1.1 X 10-,<) gives it such low solubility' that it can be safely consumed. What is the molar solubility' of BaSO4. What is its solubility' in grams per 100 g of water?