A researcher put 10.0 moles of N2O into a 2-L container at some temperature where it decomposes according to the following: N2O = 2N2+ O2. At equilibrium 2.20 moles of N2O remain. Calculate the Kc for the reaction.
How am I supposed to fill in the ICE chart if Kc isn't given for me to plug into the constant equation?
C'mon. It takes us a little time to type this stuff in. I've posted a response to your earlier post.
To calculate Kc for the reaction, you need the initial and equilibrium concentrations of the reactants and products. However, in this case, Kc is not provided.
Since the problem states that there were initially 10.0 moles of N2O in a 2-L container, we can calculate the initial concentration of N2O using the formula c = n/V, where c is the concentration, n is the number of moles, and V is the volume:
Initial concentration of N2O = (moles of N2O) / (volume of container)
= 10.0 moles / 2 L
= 5.0 M
At equilibrium, 2.20 moles of N2O remain. This means the reaction has consumed 10.0 - 2.20 = 7.80 moles of N2O.
To complete the ICE chart, start with the initial concentrations, set up a stoichiometry table, and deduce the equilibrium concentrations. Assume x moles of N2O react. The change in moles for N2O is -x, and the changes in moles for N2 and O2 are 2x and x, respectively.
ICE Chart:
N2O 5.0 M - x
N2 2x
O2 x
Let's focus on N2O. At equilibrium, the concentration of N2O is given as 2.20 moles. So, the equilibrium concentration of N2O is 2.20 M.
Now, rewrite the ICE chart using the given equilibrium values:
N2O 2.20 M
N2 2x
O2 x
Now, let's substitute the values obtained into the expression for Kc:
Kc = ([N2]^2 * [O2]) / [N2O]^2
Kc = (2x)^2 * x / (2.20)^2
Simplifying, we get:
Kc = 4x^3 / (2.20)^2
Since we don't know the value of x, we can't calculate the exact value of Kc. However, you can solve for x by assuming it to be small and plug that value into the expression if you want to obtain an approximate value for Kc.
To calculate the equilibrium constant, Kc, for the reaction N2O = 2N2 + O2, you need the concentrations (or partial pressures) of the reactants and products at equilibrium. The given information tells us that initially, 10.0 moles of N2O were placed in a 2-L container, and at equilibrium, only 2.20 moles of N2O remain.
Since the reaction is given as a decomposition reaction, it means that the concentration of N2O decreases as the reaction proceeds and the concentrations of N2 and O2 increase. To fill in the ICE (Initial, Change, Equilibrium) chart, follow the steps below:
1. Start with the initial conditions:
N2O: 10.0 moles (initial) - x (change) = 10.0 - x
N2: 0 moles (initial) + 2x (change) = 2x
O2: 0 moles (initial) + x (change) = x
2. Determine the "change" in moles for each species based on the stoichiometry of the reaction. Since 1 mole of N2O is converted to 2 moles of N2 and 1 mole of O2, we can use the value of x to determine the moles of N2 and O2 produced.
3. Set up an expression for the equilibrium concentration of N2O, using the given value of 2.20 moles remaining at equilibrium:
N2O = 10.0 - x = 2.20
4. Solve for x to find the value of x, which represents the "change" in moles.
5. Once you have obtained the value of x, substitute it back into the expressions for the equilibrium concentrations of N2 and O2:
N2 = 2x
O2 = x
6. Finally, use the equilibrium concentrations to calculate the equilibrium constant, Kc, using the formula:
Kc = [N2]^2 * [O2] / [N2O]
Substitute the equilibrium concentrations you have obtained for [N2], [O2], and [N2O] into the Kc expression, and perform the necessary calculations to determine the value of Kc.
Note: It's important to make sure that the units used for concentration are consistent (either moles per liter or partial pressures) when calculating Kc.