These Are A Portion Of The Questions That Are Needed To Be A ✓ Solved

These Are A Portion Of the Questions That Are Needed To Be Answered I

These are a portion of the questions that are needed to be answered. If you decided to complete this assignment I will screenshot and send the remaining 9 questions.

1. The equilibrium constant KP for the reaction is 203 at a certain temperature. Calculate PO2 if P NO = 0.411 atm and P NO2 = 0.385 atm.

2. For the reaction 2NO2(g) ⇌ 2NO(g) + O2(g), given P O2 = atm. Enter your answer in the provided box.

3. For the reaction N2(g) + 3H2(g) ⇌ 2NH3(g), KP is 0.000817 at 683°C. What is Kc for the reaction? The answer is 3.

4. A reaction vessel contains NH3, N2, and H2 at equilibrium at a certain temperature. The equilibrium concentrations are [N2] = 0.41 M, [H2] = 1.26 M, and [NH3] = 0.40 M. Calculate the equilibrium constant, Kc, if the reaction is represented as (1/2) N2(g) + (3/2) H2(g) ⇌ NH3(g).

5. Ammonium carbamate, NH4CO2NH2, decomposes as follows: NH4CO2NH2(s) ⇌ 2NH3(g) + CO2(g). Starting with only the solid, it is found that at 40°C the total gas pressure (NH3 and CO2) is 0.263 atm. Calculate the equilibrium KP.

6. A 2.180 mole quantity of NOCl was initially placed in a 1.150 L reaction chamber at 400°C. After equilibrium was established, it was found that 28.60 percent of the NOCl has dissociated: 2NOCl(g) ⇌ 2NO(g) + Cl2(g). Calculate the equilibrium constant KC for the reaction.

7. At 1000 K, a sample of pure NO2 gas decomposes: 2NO2(g) ⇌ 2NO(g) + O2(g). The equilibrium constant KP is 158. Analysis shows that the partial pressure of O2 is 0.43 atm at equilibrium. Calculate the pressure of NO and NO2 in the mixture.

Sample Paper For Above instruction

Introduction

Understanding equilibrium constants is fundamental in chemical thermodynamics and reaction kinetics. These constants, KP and KC, provide insights into the extent of chemical reactions under various conditions. This paper addresses six specific problems related to equilibrium constants, covering calculations for partial pressures and concentrations, as well as the application of the equilibrium constant definitions in different scenarios. Accurate determination of these values is crucial for controlling chemical processes in industrial and laboratory settings.

Problem 1: Calculating PO2 from KP and partial pressures of NO and NO2

The first problem involves the equilibrium constant KP for a reaction at a known temperature, with known partial pressures of NO and NO2. The reaction is typically represented as:

2NO2(g) ⇌ 2NO(g) + O2(g)

The equilibrium expression for KP is:

KP = P_NO² * P_O2 / P_NO2²

Given KP = 203, P_NO = 0.411 atm, and P_NO2 = 0.385 atm, we can calculate P_O2:

P_O2 = KP * P_NO2² / P_NO²

Substituting the known values:

P_O2 = 203 (0.385)² / (0.411)² ≈ 203 0.148 / 0.169 ≈ 203 * 0.876 ≈ 178. adherent to the calculation and the provided constants.

Problem 2: Determining P O2 in Nitrogen and Hydrogen Reaction

In a scenario where P_O2 is given (though unspecified in the fragment), the goal is to compute the partial pressure based on established equilibrium relations. The reaction involves nitrogen and hydrogen gases reacting to form ammonia, with a known equilibrium constant KP of 0.000817 at 683°C. The relation between KP and KC can be derived considering the ideal gas law and the reaction's equilibrium expression.

The conversion between KP and KC involves the temperature-dependent pressure-volume relationship. Using the relation:

KC = KP * (RT)^Δn

where Δn is the change in moles of gas between reactants and products, R is the gas constant, and T is temperature in Kelvin, calculations proceed after substituting known values.

Problem 3: Calculating the Equilibrium Constant KC from Concentrations

Given molar concentrations of N₂, H₂, and NH₃ at equilibrium, the reformulated reaction is:

(1/2) N₂(g) + (3/2) H₂(g) ⇌ NH₃(g)

The equilibrium constant KC is expressed as:

KC = [NH₃] / ([N₂]^{1/2} * [H₂]^{3/2})

Substituting the given concentrations:

KC = 0.40 / (0.41^{1/2} 1.26^{3/2}) ≈ 0.40 / (0.640 1.413) ≈ 0.40 / 0.905 ≈ 0.442

The calculation yields a value around 0.44, indicating the extent of the reaction at given conditions.

Problem 4: Determining the Equilibrium Constant KP from Gas Pressures

The decomposition of ammonium carbamate results in gaseous products NH3 and CO2. With total pressure at 0.263 atm at 40°C, and the decomposition reaction:

NH₄CO₂NH₂(s) ⇌ 2NH₃(g) + CO₂(g)

The equilibrium expression for KP is:

KP = (P_NH3)^2 * P_CO2

Assuming ideal gas behavior, the partial pressures of NH₃ and CO₂ sum to the total pressure, and their ratio can be derived from the stoichiometry. Let x be the partial pressure of NH₃, then P_CO2 = (Total pressure) - 2x (since for every 2 mols of NH₃, 1 mol of CO₂ is produced). With additional data or assumptions, KP can be computed accordingly.

Problem 5: Calculating KC from Dissociation Data of NOCl

Initial moles = 2.180 mol in 1.150 L at 400°C. The percent dissociation is 28.60%. The molar amounts at equilibrium can be calculated:

Disassociated moles = 2.180 * 0.286 ≈ 0.623 mol

Remaining NOCl = 2.180 - 0.623 ≈ 1.557 mol

Partial pressures from ideal gas law:

P = (nRT)/V

Using the partial pressures to compute KC:

KC = [NO]^2 * [Cl₂] / [NOCl]²

Expressed numerically, following the detailed stoichiometry, yields the equilibrium constant.

Problem 6: Calculating Nitrogen Oxide Pressures at Equilibrium

The decomposition of NO₂ at 1000 K with KP = 158 and partial pressure of O₂ at 0.43 atm involves solving the equilibrium expression:

KP = (P_NO)^2 * P_O2 / (P_NO₂)^2

Given P_O2 = 0.43 atm, solving for P_NO and P_NO₂ involves algebraic manipulation, assuming initial moles or using the partial pressures to determine the equilibrium concentrations. Numeric solutions will ensure the coherence of the partial pressures with the equilibrium constant.

Conclusion

These problems demonstrate the application of equilibrium expressions, the relationship between KP and KC, and the use of partial pressures and concentrations to determine the extent of chemical reactions. Mastery of these calculations is essential for chemists analyzing reaction systems, optimizing industrial reactions, and predicting product yields under various conditions.

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