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3 years ago

Which of the following is not true for an exothermic reaction?

Chemistry

1 answer:

Gennadij [26K]3 years ago

4 0

Answer:

The false statement regarding an exothermic reaction is: A. the products have a higher enthalpy than reactants

Explanation:

An exothermic reaction is a type of chemical reaction that involves the release of energy from the system to the surroundings. Thus increasing the temperature of the surroundings.

In this reaction, the enthalpy or energy of the reactants is greater than the enthalpy or energy of the products. ( $\Delta H_{f} (Products) < \Delta H_{f} (reactants)$ )

As the enthalpy change of a reaction: $\Delta H_{r} = \sum \Delta H_{f} (Products) - \sum \Delta H_{f} (reactants)$

Therefore, the enthalpy change for an exothermic reaction is negative ( $\Delta H_{r} < 0$ )

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Given that Δ H ∘ f [ Br ( g ) ] = 111.9 kJ ⋅ mol − 1 Δ H ∘ f [ C ( g ) ] = 716.7 kJ ⋅ mol − 1 Δ H ∘ f [ CBr 4 ( g ) ] = 29.4 kJ

JulsSmile [24]

Answer:

283.725 kJ ⋅ mol − 1

Explanation:

C(s) + 2Br2(g) ⇒ CBr4(g) , Δ H ∘ = 29.4 kJ ⋅ mol − 1

$\frac{1}{2}$ Br2(g) ⇒ Br(g) , Δ H ∘ = 111.9 kJ ⋅ mol − 1

C(s) ⇒ C(g) , Δ H ∘ = 716.7 kJ ⋅ mol − 1

4*eqn(2) + eqn(3) ⇒ 2Br2(g) + C(s) ⇒ 4 Br(g) + C(g) , Δ H ∘ = 1164.3 kJ ⋅ mol − 1

eqn(1) - eqn(4) ⇒ 4 Br(g) + C(g) ⇒ CBr4(g) , Δ H ∘ = -1134.9 kJ ⋅ mol − 1

so,

average bond enthalpy is $\frac{1134.9}{4}$ = 283.725 kJ ⋅ mol − 1

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3 years ago

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Answer:

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Explanation:

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3 years ago