## Chemistry (4th Edition)

Published by McGraw-Hill Publishing Company

# Chapter 15 - Questions and Problems - Page 710: 15.49

#### Answer

The reverse reaction will occur, so the forward reaction will not occur.

#### Work Step by Step

1. Write the equilibrium constant expression: - The exponent of each concentration is equal to its balance coefficient. $$K_C = \frac{[Products]}{[Reactants]} = \frac{[ H_2O ][ CO ]}{[ H_2 ][ CO_2 ]}$$ 2. At equilibrium, these are the concentrations of each compound: $[ H_2 ] = 3.0 \times 10^{-5} \space M - x$ $[ CO_2 ] = 3.3 \times 10^{-4} \space M - x$ $[ H_2O ] = 1.6 \times 10^{-2} \space M + x$ $[ CO ] = 6.25 \times 10^{-3} \space M + x$ $$1.11 = \frac{(1.6 \times 10^{-2} + x)(6.25 \times 10^{-3} + x)}{(3.0 \times 10^{-5} - x)(3.3 \times 10^{-4} - x)}$$ $x_1 = -0.0043$ $x_2 = 0.21$ $x_2$ is invalid, because $[H_2] = 3.0 \times 10^{-5} - 0.21 =$ negative number. A concentration cannot be negative. So x = -0.0043 Which means that the concentration of reactants will increase, and the same of the products will decrease, which means that the reverse reaction will occur.

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