Consider the reaction: 2 SO2 (g) + O2 (g) ⇋ 2 SO3 (g)
4.0 mol of SO2 and 2.2 mol of O2 were placed in a 10 L flask and allowed to reach equilibrium. The flask was found to contain 2.6 mol SO2. Calculate the equilibrium constant K for this reaction.
(SO2) = mols/L = 4.0 mols/10 L = 0.40 M
(O2) = 2.2/10 = 0.22 M
........................ 2SO2 (g) + O2 (g) ⇋ 2SO3 (g)
I.......................0.40M..........0.22 M...........0
C.........................-2x...........-x...........+2x
E......................0.40 - 2x.....0.22-x........2x
The problem tells you that mols SO2 at equilibrium was 2.6 mols. That in 10 L is 2.6/10 = 0.26 M so you know that 0.40-2x = 0.26. Solve for x, evaluate 0.40-2x = (SO2) at equilibrium, 0.22-x = (O2) at equilibrium, and 2x = (SO3) at equilibrium. Substitute those evaluations into Keq expression and solve for K. Post your work if you have any questions.
Okay thanks that makes sense but I am confused for this next one
Consider the reaction: 2 IF5 (g) + I4F2 (g) ⇋ 3 I2 (g) + 6 F2 (g)
6.0 mol of IF5 and 8.0 mol of I4F2 are placed in a 5.0 L container. At equilibrium, 6.0 mol of I4F2 are present. Calculate K for this reaction.
Well, let's solve this problem step by step! First, we need to find the moles of O2 remaining at equilibrium. We started with 2.2 mol and ended up with 2.6 mol of SO2, so the difference is (2.2 - 2.6) = -0.4 mol.
Now, if we look at the balanced equation, we can see that for every 2 moles of SO2 consumed, 1 mole of O2 is consumed. So, since we have -0.4 mol of O2 left, we can say that (2 * -0.4) = -0.8 mol of SO2 has been consumed.
Now we can calculate the moles of SO2 remaining. We had 4.0 mol initially, and we consumed -0.8 mol. Therefore, the remaining moles of SO2 are (4.0 - (-0.8)) = 4.8 mol.
Finally, we can calculate the equilibrium constant (K) using the formula:
K = [SO3]^2 / ([SO2]^2 * [O2])
Substituting the values we have:
K = (2.6^2) / (4.8^2 * 0.4)
After some calculations, we get K ≈ 0.151
So, the equilibrium constant (K) for this reaction is approximately 0.151.
To calculate the equilibrium constant (K) for the reaction, we need to know the concentrations (or pressures) of each species at equilibrium. Since we are given the number of moles of SO2 at equilibrium, we need to calculate the concentrations of SO2, O2, and SO3.
We can use the information provided to set up an ICE (Initial, Change, Equilibrium) table as follows:
Reaction: 2 SO2 (g) + O2 (g) ⇋ 2 SO3 (g)
_____________________________________________
| SO2 | O2 | SO3
_____________________________________________
Initial | 4.0 mol | 2.2 mol | 0 mol
Change | -2.6 mol | -2.6 mol |+2.6 mol
Equilibrium | 1.4 mol | -0.4 mol | 2.6 mol
Using the ideal gas law, we can calculate the concentrations of SO2, O2, and SO3 at equilibrium:
[SO2] = n / V = 1.4 mol / 10 L = 0.14 M
[O2] = n / V = (-0.4 mol) / 10 L = -0.04 M (Note: Negative concentration means that O2 is consumed and its concentration is zero at equilibrium)
[SO3] = n / V = 2.6 mol / 10 L = 0.26 M
Now, we can substitute the concentrations into the equilibrium expression:
K = ([SO3]^2) / ([SO2]^2 * [O2])
K = (0.26^2) / (0.14^2 * 0.04)
K = 1.7676 / 0.000784
K ≈ 2254.8
Therefore, the equilibrium constant (K) for this reaction is approximately 2254.8.
To calculate the equilibrium constant (K) for this reaction, we need to use the equation:
K = ([SO3]^2) / ([SO2]^2 * [O2])
Given that the initial amounts of SO2, O2, and the amount of SO2 at equilibrium are 4.0 mol, 2.2 mol, and 2.6 mol respectively, we can substitute these values into the equation.
1. Calculate the concentration of SO3 at equilibrium:
[SO3] = (Amount of SO2 at equilibrium) / (Volume of flask)
[SO3] = 2.6 mol / 10 L
[SO3] = 0.26 M
2. Calculate the concentration of SO2 and O2 at equilibrium:
[SO2] = (Amount of SO2 at equilibrium) / (Volume of flask)
[SO2] = 2.6 mol / 10 L
[SO2] = 0.26 M
[O2] = (Amount of O2 initially - Amount of O2 reacted) / (Volume of flask)
[O2] = (2.2 mol - 0 mol) / 10 L
[O2] = 0.22 M
3. Substitute the concentrations into the equation for K:
K = ([SO3]^2) / ([SO2]^2 * [O2])
K = (0.26 M)^2 / (0.26 M)^2 * (0.22 M)
K = 0.0676 / (0.0676 * 0.22)
K = 0.0676 / 0.014872
K = 4.54
Therefore, the equilibrium constant (K) for this reaction is 4.54.