1. Consider the set of liquid-phase series reactions 2A3B →2C. The first reaction is secondorder, while the second reaction is first order. The rate constants at 40°C are k₁ = 0.3 L/mol/hand k₂ = 1 h¹, while EA₁ = 7 kcal/mol and EA2 = 0.20 kcal/mol. CA0 = 8 mol/L.a) Plot the concentrations of each species in a 100 L batch reactor at 40°C and 100°C over10 hours (assume the reaction is liquid phase at both temperatures).b) What is the maximum concentration of species B at both temperatures?c) At what time do you get the maximum concentration of B at both temperatures?

To address the problem, we need to follow these steps:1. Determine the reaction rate constants at…

Answered: 1. Consider the set of liquid-phase…

Introduction to Chemical Engineering Thermodynamics

8th Edition

ISBN:9781259696527

Author:J.M. Smith Termodinamica en ingenieria quimica, Hendrick C Van Ness, Michael Abbott, Mark Swihart

Publisher:J.M. Smith Termodinamica en ingenieria quimica, Hendrick C Van Ness, Michael Abbott, Mark Swihart

Chapter1: Introduction

Section: Chapter Questions

Problem 1.1P

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A set of four reversible elementary reactions that describe the oxidation of CO to CO2 in the presence of water is given below. We will discuss these further when we get to chemical kinetics. Here you are to consider the thermodynamics of these reactions. a. From the property tables in your textbook, compile a table of the following quantities for each of the four reactions: the enthalpy of reaction for the forward reaction at Tref (AHR at 298 K), the partial-pressure equilibrium constant at 1800 K (Kp at 1800 K), and the partial-pressure equilibrium constant at 2000 K (Kp at 2000 K). b. In which direction (left-to-right, right-to-left, or neither) is each of the four reactions exothermic? c. In which direction (toward the left, toward the right, or neither) will the equilibrium of each reaction shift if the pressure of the system is increased from 1 atm to 5 atm, while the temperature remains constant?
Consider the following reaction: 4 NH3(g) + 5 O2(g) = 6 H20(g) + 4 NO(g) Calculate AG at 547 K for a mixture of 2.3 bar NH3(g), 3.0 bar O2(g), 4.9 bar H20(g), and 3.6 bar NO(g). State whether or not the reaction is spontaneous at these conditions. **AG° values can be found in textbook appendix** O -929 kJ/mol ; spontaneous -929 kJ/mol ; nonspontaneous O 929 kJ/mol ; spontaneous O 929 kJ/mol ; nonspontaneous
Question 3 A vapour-liquid mixture of water (1) and furfural CSH4O2 (2) is maintained at 1.013 bar and 109.5 °C. Under these conditions, it is found that x1 = 0.10 and y1 = 0.81. The temperature of the mixture %3D is changed to 100.6 °C, and some (but not all) of the vapour condenses. The vapour phase may be assumed to behave as an ideal gas. The liquid-phase activity coefficients are independent of temperature but vary with concentration according to the van Laar equations. Determine: (a) The activity coefficients for water and furfural at 1.013 bar and 109.5 °C [9%] (b) The van Laar constants [6% ] Vapour pressure data: T=1006°C, Pa =1.0352 bar, Pl= 0.1193 bar T=109.5°C, Psal -14088 bar, Pal -0.1690 bar %3D
For the gas-phase reaction A = 2B operated at 1.0 bar: Initially, the reactor contains only A, at 1.0 mol. a. Prove that the reaction coordinate (moles of A consumed) equals the following: (see attactched image) b. Make a plot of the reaction coordinate versus temperature between 200 and 400 K for exothermic reaction conditions, making use of the van’t Hoff Equation. Data: Ka = 5.0 at 298 K, ΔHrxn0 = -50.0 kJ mol-1. Neglect the ΔCp term. c. Repeat Part b for endothermic conditions. Ka = 5.0 at 298 K, ΔHrxn0 = +50.0 kJ mol-1. Neglect the ΔCp term. d. If the reaction described in part b (i.e., exothermic conditions) were operated in a simple batch reactor, would the reaction mixture’s temperature increase, decrease, or remain constant as the reaction progressed? To maximize conversion, would you recommend that heat be added to, or removed from, the reactor?
Assuming the reaction below is at equilibrium, which of the following changes will drive the reaction to the left? C(s) + O2(g)= CO2(g) AH = -393.5 kJ/mol Increasing the temperature. Adding O2(g). 1. 2. 3. Removing C(s). 1 only 2 only 3 only 1 and 3 а. b. с. d. е. 2 and 3
Why is the reaction with a negative Gibbs free energy always spontaneous? a) more energy is needed to form the products than the reactants, so energy must be added b) less energy is needed to form the products than the amount of energy present in the reactants c) a negative deltaG corresponds to a negative deltaE, which is always spontaneous d) a negative Gibbs free energy, deltaG, always corresponds with a postive entropy, deltaS.
Calculate the standard free energy for the reaction at 312.0°C. (you may use any valid technique - there may be more than one way to do a question like this): Al2O3(s)) + 3 H2(g) → 2 Al(s) + 3 H2O(g)

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