Question

Consider the following intermediate chemical equations.(IMAGE) -205.7 kJ -113.4 kJ -14.3 kJ 78.0 kJ

Answer 1

Approximately
`-205.7\; \rm kJ`.

Step-by-step explanation:

This question can be solved using Hess's Law.

Start by considering: how can the first three reactions (with known
`\Delta H` values) be combined to produce the reaction
`\rm CH_4\; (g) + 4\; \rm Cl_2\; (g) \to CCl_4\; (g) + 4\; HCl\; (g)`?

Here's one possible combination:

• Include the first reaction once, without inverting.
• Invert the second reaction and include it once.
• Include the third reaction after multiplying all its coefficients by two.

In other words, if
`(1)`,
`(2)`, and
`(3)` denote the three reactions with know
`\Delta H` values, respectively, then
`1 * (1) - 1 * (2) + 2* (3)` will give the required reaction
`\rm CH_4\; (g) + 4\; \rm Cl_2\; (g) \to CCl_4\; (g) + 4\; HCl\; (g)`.

By Hess's Law, the
`\Delta H` value of the reaction
`\rm CH_4\; (g) + 4\; \rm Cl_2\; (g) \to CCl_4\; (g) + 4\; HCl\; (g)` will thus be:

`\begin{aligned}&1 * \Delta H_1 - 1* \Delta H_2 + 2* \Delta H_3\\ &= 1 * 74.6\; \rm kJ - 1 * 95.7\; \rm kJ +2 * (-92.3\; \rm kJ) \\ &= -205.7\; \rm kJ\end{aligned}`.