At elevated temperatures, SbCl5 gas decomposes into SbCl3 gas and Cl2 gas as shown by the following equation: SbCl5(g)<=> SbCl3(g) + Cl2(g)
1. What is the concentration in moles per liter of SbCl5 in the container before any decomposition occurs?
2. What is the pressure in atmospheres of SbCl5 in the container before any decomposition occurs?
b)If the SbCl5 is 29.2 percent decomposed when equilibrium is
established at 182°C, calculate the values for equilibrium constants Kp and Kc, for this decomposition reaction.
Kc=?
Kp=?
c.In order to produce some SbCl5, a 1.00 mole sample of SbCl3 is first
placed in an empty 2.00 liter container maintained at a temperature
different from 182ºC. At this temperature, Kc, equals 0.117. How many moles of Cl2 must be added to this container to reduce the number of
moles of SbCl3 to 0.500 mole at equilibrium?
I think you need more information. Do you have a number of moles? a volume? pressure?
1. (Cl2) = 0; (SbCl3) = 0; (SbCl5) = whatever you started with.
To answer these questions, we need to use the principles of equilibrium and the expression of the equilibrium constant. The equilibrium constant is a quantitative measure of the extent of a reaction at equilibrium.
1. To determine the initial concentration of SbCl5 before any decomposition occurs, we assume that no decomposition has taken place. Therefore, the concentration of SbCl5 is equal to its initial concentration. Without additional information, we cannot determine the concentration of SbCl5 in moles per liter.
2. To calculate the pressure in atmospheres of SbCl5 in the container before any decomposition occurs, we need to know the conditions under which the pressure is measured. If the pressure is measured at standard conditions (25°C and 1 atm), we can use the ideal gas law to calculate the pressure:
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. Rearranging the equation, we can solve for the number of moles:
n = PV / RT
If the volume and temperature are known, we can substitute the values into the equation to calculate the number of moles of SbCl5. Then, we can convert the moles to a concentration in moles per liter.
b) To calculate the equilibrium constant for the decomposition reaction given the percent decomposed and the temperature, we need to use the following equation:
Kp = (P(SbCl3) × P(Cl2)) / (P(SbCl5)) (1)
Kc = ([SbCl3] × [Cl2]) / [SbCl5] (2)
We can use the given percent decomposition to calculate the concentrations/pressures of each species at equilibrium:
P(SbCl5)eq = P(SbCl5)initial × (1 - percent decomposition)
P(SbCl3)eq = P(SbCl5)initial × percent decomposition
P(Cl2)eq = P(SbCl5)initial × percent decomposition
Using these equilibrium pressures and the ideal gas law, we can calculate the equilibrium concentrations of each species in moles per liter. Then, substitute these concentrations into equations (1) and (2) to calculate the equilibrium constants Kp and Kc.
c) The equilibrium constant Kc relates the moles of reactants and products at equilibrium. Based on the equation for the decomposition reaction:
SbCl5(g) <=> SbCl3(g) + Cl2(g)
The stoichiometry of the reaction is 1:1:1. This means that for every mole of SbCl5 that decomposes, we produce 1 mole of SbCl3 and 1 mole of Cl2. Given the equilibrium constant Kc, we can set up an expression to determine the moles of Cl2 needed to achieve a certain number of moles of SbCl3 at equilibrium:
Kc = ([SbCl3] × [Cl2]) / [SbCl5]
Since the initial concentration of SbCl3 is given as 1.00 mole and we want to reduce it to 0.500 mole at equilibrium, we can rearrange the equation as follows:
[SbCl5] = ([SbCl3] × [Cl2]) / Kc
Then substitute the given values for [SbCl3] and Kc, and solve for [Cl2] to find the number of moles of Cl2 that need to be added to the container.