We can say that this reaction is (endothermic or exothermic) and that the sign of $$ \Delta \mathrm{H} $$ is (positive or negative). The energy of products is (more or less) than energy of reactants. $$ 2 \mathrm{C}_{6} \mathrm{H}_{6}(\mathrm{l})+15 \mathrm{O}_{2}(\mathrm{~g}) \rightarrow 12 \mathrm{CO}_{2}(\mathrm{~g})+6 \mathrm{H}_{2} \mathrm{O}(\mathrm{g})+6278 \mathrm{~kJ} $$

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To determine whether the reaction is endothermic or exothermic, we can analyze the given information about the reaction:

$$
2 \mathrm{C}_{6} \mathrm{H}_{6}(\mathrm{l}) + 15 \mathrm{O}_{2}(\mathrm{g}) \rightarrow 12 \mathrm{CO}_{2}(\mathrm{g}) + 6 \mathrm{H}_{2} \mathrm{O}(\mathrm{g}) + 6278 \mathrm{~kJ}
$$

1. **Exothermic or Endothermic**: The reaction releases energy, as indicated by the positive value of $$ 6278 \, \text{kJ} $$ on the product side. This means that energy is released to the surroundings during the reaction, which characterizes it as **exothermic**.

2. **Sign of $$ \Delta H $$**: Since the reaction releases energy, the change in enthalpy ($$ \Delta H $$) is negative. Therefore, we can say that the sign of $$ \Delta H $$ is **negative**.

3. **Energy of Products vs. Reactants**: In an exothermic reaction, the energy of the products is **less** than the energy of the reactants because energy is released during the reaction.

In summary:
- The reaction is **exothermic**.
- The sign of $$ \Delta H $$ is **negative**.
- The energy of products is **less** than the energy of reactants.
The reaction is exothermic, $$ \Delta H $$ is negative, and the energy of products is less than the energy of reactants.

We can say that this reaction is (endothermic or exothermic) and that the sign of
is (positive or negative). The energy of products is (more or less) than energy of
reactants.

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We can say that this reaction is (endothermic or exothermic) and that the sign of is (positive or negative). The energy of products is (more or less) than energy of reactants.