$$\ce{S(s) + O2(g) -> SO2(g)}$$

The reaction of sulfur with oxygen is written in equation form above. This equation can be interpreted in all of the following ways except
(A) either $\ce{S(s)}$ or $\ce{O2(g)}$ will be completely used up
(B) $Q$ must be close to $1.0$ since there is one mole of gas on each side of the equation
(C) this reaction goes to completion
(D) adding $\ce{O2}$ will change the equilibrium constant

I do not know how to extrapolate $Q$ from just an equation with no other data. However, I also learned that the equilibrium constant is, well, constant, and adding O$_2$ should change the equilibrium concentrations, not the equilibrium constant.

The correct answer is (B). How would (D) be a valid interpretation?

  • $\begingroup$ I'd say it wouldn't. Your interpretation is correct, and the test is inaccurate. $\endgroup$ May 1 '19 at 6:49

D is an invalid statement because the ''equilibrium constant'' for a certain reaction is constant as long as the temperature of the system is constant. When the concentration of a species Is changed according to Le Chatelier's principle the equilibrium system will change it's rates of the forward and backward reactions so as to balance the concentrations thus keeping the equillibrium constant the same.


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