A mixture containing 0.5 mol of H2 (g) and 0.5 mol of I2 (g) was introduced into a 10L container at a temperature of 448 C. At this temperature the equilibrium constant Kp is 50.
i. What is the total pressure in the container before the reaction starts and at equilibrium.
ii. How many moles of hydrogen remain at equilibrium
iii. What is the partial pressure of each component in the equilibrium mixture
Thanks!!!
I. Use PV = nRT for total P. I think that is about 5.9 atm but you need to do it more carefully than that.
........H2 + I2 ==> 2HI
I......2.95...2.95...0
C......-p.....-p....+2p
E.....2.95-p..2.95-p..2p
III.
Substitute the equil line into Kp and solve for p, then 2p, then 2.95-p.
II. Use PV = nRT; substitute p for H2 and solve for n for mols H2.
To solve this problem, we need to understand the concept of equilibrium and how to calculate the total pressure and partial pressures in a gaseous mixture.
1. Total Pressure:
The total pressure in the container before the reaction starts can be calculated using the Ideal Gas Law equation: PV = nRT, where P is the pressure, V is the volume, n is the number of moles, R is the ideal gas constant, and T is the temperature in Kelvin.
Before the reaction starts, both hydrogen (H2) and iodine (I2) are in their gaseous states and have equal number of moles (0.5 mol each). So the total pressure can be calculated by summing up the partial pressures of each gas using Dalton's Law of Partial Pressures:
P(total) = P(H2) + P(I2)
2. Equilibrium:
At equilibrium, the reaction has reached a balance between the forward and reverse reactions. The equilibrium constant, Kp, expresses this balance and is given by the ratio of the partial pressures of the products to the reactants, each raised to the power of their respective stoichiometric coefficient.
Kp = (P(H2)^x * P(I2)^y) / (P(H2O)^a * P(I)^b)
In this case, as the equation is not given, we cannot fully determine the values of x, y, a, and b. However, we can calculate the equilibrium partial pressures using this equation and substitute any known values.
3. Partial Pressures:
To calculate the partial pressures at equilibrium, we can assume that the reaction has reached equilibrium and denote the number of moles of H2 remaining as n(H2).
From the balanced chemical equation, assuming the reaction is:
H2 (g) + I2 (g) ↔ 2HI (g)
The stoichiometry states that for every mole of H2 that reacts, 2 moles of HI are produced. Thus, the number of moles of HI produced is 2 * n(H2).
Using this information, we can write the expressions for partial pressures at equilibrium:
P(H2) = (0.5 - n(H2)) * R * T / V
P(I2) = (0.5 - n(H2)) * R * T / V
P(HI) = 2 * n(H2) * R * T / V
Notice that the total pressure calculated in the first step will be equal to the sum of the partial pressures in equilibrium (P(total) = P(H2) + P(I2) + P(HI)). You can use this equation to find the total pressure after the reaction has reached equilibrium.
To find the remaining moles of hydrogen at equilibrium, you need to know more about the equilibrium constant or perform additional calculations regarding the stoichiometry of the reaction.
In summary:
i. The total pressure in the container before the reaction starts can be calculated using the Ideal Gas Law. At equilibrium, the total pressure can be found by summing up the partial pressures of each gas.
ii. The number of moles of hydrogen remaining at equilibrium depends on the stoichiometry of the reaction and the equilibrium constant. More information is needed to determine the moles of hydrogen remaining.
iii. The partial pressures at equilibrium can be calculated using the equilibrium expression and the known values of the number of moles of hydrogen remaining.