Consider the reaction:
CO(g) + H2O (g) <---> CO2(g)+H2(g)
which of the following initial composition will NOT achieve equlibrium?
a) 0.25 mol H2, 0.35 mol CO2, and 0.25 mol CO
b) 0.25 mol CO, 0.75 mol H2O
c) 1.0 mol of CO2 and 0.50 mol H2
d) 0.85 mol CO and 0.25 mol H2
how do i do this? i can't make an ICE table because no volume is given. which means i can't find Keq or Q . So how else am i supposed to reason this?
I would look at it this way.
A is missing only H2O so it can move to the left to reach equilibrium.
B is missing the right side so it can move to the right.
C is missing the left side so it can move to the left.
D is missing something on the left AND something on the right. That means CO must decrease and H2 must decrease. If it moves to the left to decrease H2 what happens to the CO2. It increases but it must decrease. Right? If it moves to the right to decrease CO2 what happens to the H2. It increases but it must decrease. So this one can't work.
yes that makes sense, thank you!!
To determine which initial composition will not achieve equilibrium, we can reason by considering the reaction stoichiometry and Le Chatelier's principle.
Let's analyze each option:
a) 0.25 mol H2, 0.35 mol CO2, and 0.25 mol CO:
From the coefficients of the balanced reaction, we can see that 1 mol of CO reacts with 1 mol of H2O, forming 1 mol of CO2 and 1 mol of H2. Therefore, the initial composition in option a) has more H2 and CO2 than CO, which means the reaction can proceed to reach equilibrium. So, option a) is not the correct answer.
b) 0.25 mol CO, 0.75 mol H2O:
From the balanced reaction, we can see that 1 mol of CO reacts with 1 mol of H2O, forming 1 mol of CO2 and 1 mol of H2. In option b), there is an excess of H2O compared to CO, which could potentially inhibit the reaction from reaching equilibrium. So, option b) is a candidate for the correct answer.
c) 1.0 mol of CO2 and 0.50 mol H2:
In this case, the initial composition has more CO2 than H2, which suggests that the reaction can proceed to reach equilibrium. So, option c) is not the correct answer.
d) 0.85 mol CO and 0.25 mol H2:
From the balanced reaction, we can see that 1 mol of CO reacts with 1 mol of H2O, forming 1 mol of CO2 and 1 mol of H2. In option d), there is an excess of CO compared to H2, which could potentially inhibit the reaction from reaching equilibrium. So, option d) is a candidate for the correct answer.
Based on our reasoning, the correct answer would be either option b) or option d). To make a final decision, you may need additional information or experimental data.
To determine which of the given compositions will NOT achieve equilibrium, we can analyze the reaction and apply the concept of Le Chatelier's Principle. Le Chatelier's Principle states that if a system at equilibrium is subjected to a change, it will adjust to minimize the effect of that change.
In this case, let's analyze each composition:
a) 0.25 mol H2, 0.35 mol CO2, and 0.25 mol CO:
To achieve equilibrium, the reaction needs CO and H2O to produce more CO2 and H2. Since there is an excess of CO2 and a limited amount of CO and H2O, this composition will not achieve equilibrium.
b) 0.25 mol CO, 0.75 mol H2O:
The reaction needs CO2 and H2, but there is no CO2 present. Therefore, this composition will not achieve equilibrium.
c) 1.0 mol of CO2 and 0.50 mol H2:
This composition has the necessary components, CO2 and H2, to reach equilibrium. It is possible for this composition to achieve equilibrium.
d) 0.85 mol CO and 0.25 mol H2:
To reach equilibrium, the reaction needs CO2 and H2. However, there is a limited amount of H2 available. Therefore, this composition will not achieve equilibrium.
In summary, the compositions that will NOT achieve equilibrium are a) 0.25 mol H2, 0.35 mol CO2, and 0.25 mol CO and d) 0.85 mol CO and 0.25 mol H2.