Calculate K p for each reaction. b. N 2 (g) + 3 H 2 (g) ⇌ 2 NH 3 (g) K c = 3.7⨉10 8 (at 298 K) c. N 2 (g) + O 2 (g) ⇌ 2 NO(g) K c = 4.10⨉10 -31 (at 298 K)

Hello everyone today, we have been given a following reaction with the K. C value at a specific temperature and asked to determine the value of K. P. Well, first we want to write the equation that we're gonna be using and that is K. P. What we want to find is equal to K. C. Times the gas constant times the temperature. And we're going to raise that to the difference in in or the difference in moles which represented by delta in. We have our k. c value which is 6.9 six times 10 to the -5. R. Gas Constant is 0.08 - one. And our temperature while in C we must convert to Kelvin by adding to 73 0.15. That gives us 333.15 Kelvin. So we're gonna write that here and then we have the difference in moles of gas and that's going to be from products to reactant. So on our product side we have one mole of ph three and one mole of BCL three which gives us two moles total. We're gonna write two minus And then we have zero moles on of gas on the reactant side. So we're going to say -0 -0 0 Calculations will yield us a final answer of 0.05, two. I hope this helped. And until next time

Ch.15 - Chemical Equilibrium

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Tro 4th Edition

Ch.15 - Chemical Equilibrium

Problem 32b,c

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Chapter 15, Problem 32b,c

Calculate K_p for each reaction.
b. N₂(g) + 3 H₂(g) ⇌ 2 NH₃(g) K_c = 3.7⨉10⁸ (at 298 K)
c. N₂(g) + O₂(g) ⇌ 2 NO(g) K_c = 4.10⨉10^-31 (at 298 K)

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Identify the relationship between Kc and Kp using the equation: Kp = Kc(RT)^(Δn), where Δn is the change in moles of gas, R is the ideal gas constant, and T is the temperature in Kelvin.

Calculate Δn for the reaction: Δn = (moles of gaseous products) - (moles of gaseous reactants). For the reaction N2(g) + 3H2(g) ⇌ 2NH3(g), Δn = 2 - (1 + 3).

Substitute the values into the equation: Kp = Kc(RT)^(Δn). Use R = 0.0821 L·atm/(mol·K) and T = 298 K.

Simplify the expression by calculating (RT)^(Δn) using the values of R, T, and Δn.

Multiply Kc by the result from the previous step to find Kp.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Equilibrium Constant (Kc and Kp)

The equilibrium constant (K) quantifies the ratio of the concentrations of products to reactants at equilibrium for a given reaction. Kc refers to concentrations in molarity, while Kp refers to partial pressures. The relationship between Kc and Kp is given by the equation Kp = Kc(RT)^(Δn), where Δn is the change in moles of gas, R is the ideal gas constant, and T is the temperature in Kelvin.

Reaction Stoichiometry

Reaction stoichiometry involves the quantitative relationships between reactants and products in a chemical reaction. In the given reaction, N2(g) + 3H2(g) ⇌ 2NH3(g), the stoichiometric coefficients indicate that one mole of nitrogen reacts with three moles of hydrogen to produce two moles of ammonia. This stoichiometry is crucial for calculating changes in concentrations or pressures at equilibrium.

Temperature Dependence of Kp

The equilibrium constant Kp is temperature-dependent, meaning its value changes with temperature. For exothermic reactions, Kp decreases with increasing temperature, while for endothermic reactions, Kp increases. Understanding this dependence is essential when calculating Kp at a specific temperature, as it directly influences the equilibrium position and the concentrations of reactants and products.

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Textbook Question

Calculate K_c for each reaction.

b. CH₄(g) + H₂O(g) ⇌ CO(g) + 3 H₂(g) K_p = 7.7x10²⁴ (at 298 K)

c. I₂(g) + Cl₂(g) ⇌ 2 ICl(g) K_p = 81.9 (at 298 K)

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Textbook Question

Coal, which is primarily carbon, can be converted to natural gas, primarily CH₄, by the exothermic reaction: C(s) + 2 H₂(g) ⇌ CH₄(g) Which disturbance will favor CH₄ at equilibrium?

a. adding more C to the reaction mixture b. adding more H₂ to the reaction mixture d. lowering the volume of the reaction mixture f. adding neon gas to the reaction mixture

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Open Question

Nitrogen dioxide dimerizes according to the reaction: 2 NO2(g) ⇌ N2O4(g) Kp = 6.7 at 298 K. A 2.25-L container contains 0.055 mol of NO2 and 0.082 mol of N2O4 at 298 K. Is the reaction at equilibrium? If not, in which direction will the reaction proceed?