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Difference between Internal Energy and Enthalpy

Which of the ...

Question

Which of the following is an exothermic reaction ?

H_{2} (g) + C l_{2 (g)} \to 2 H C l_{g}; Δ H = 184.6 k J

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N_{2 (g)} + O_{2 (g)} \to 2 N O_{(g)}; Δ H = 180.8 k J

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C_{g r a p h i t e} + H_{2} O_{(g)} \to C O_{2} + H_{2 (g)}; Δ H = - 131.4 k J

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C_{g r a p h i t e} + 2 S_{(g)} + 91.9 k J \to C S_{2 (l)}

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Solution

The correct option is C $C_{g r a p h i t e} + H_{2} O_{(g)} \to C O_{2} + H_{2 (g)}; Δ H = - 131.4 k J$
In thermodynamics, the exothermic reaction describes a reaction which releases energy from the system to its surroundings, usually in the form of heat. So for the exothermic reactions enthalpy of the reaction must be negative i.e, $Δ H$ is negative. So the correct option is C.

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Similar questions

P C l_{5} (g) + H_{2} O (g) ⟶ P O C l_{3} (g) + 2 H C l (g)

P_{w h i t e} + (3 / 2) C l_{2} (g) + \frac{1}{2} O_{2} (g) ⟶ P O C l_{3}; Δ H = - 135.5 k c a l

H_{2} (g) + C l_{2} (g) ⟶ + 2 H C l (g); Δ H = - 44.1 k c a l

P (w) + (5 / 2) C l_{2} (g) ⟶ P C l_{5} (g); Δ H = - 89.6 k c a l

H_{2} (g) + \frac{1}{2} O_{2} (g) ⟶ H_{2} O (g); Δ H = - 57.8 k c a l

N_{2} (g) + O_{2} (g) \to 2 N O (g)

Δ H = 181

2 C O_{2} (g) \to 2 C O (g) + O_{2} (g); Δ H = 566

H_{2} (g) + I_{2} (g) \to 2 H I (g); Δ H = - 94

H_{2} (g) + F_{2} (g) \to 2 H F (g); Δ H = - 541

C (g r a p h i t e) + O_{2} (g) \to C O_{2} (g); Δ H = - 393.5 k J

H_{2} (g) + \frac{1}{2} O_{2} (g) \to H_{2} O (l); Δ H = - 286.2 k J

C_{2} H_{4} (g) + 3 O_{2} (g) \to 2 C O_{2} (g) + 2 H_{2} O (l); Δ H = - 1410.8 k J

H_{2} (g) + C l_{2} (g) \to 2 H C l (g)

Δ H = - 44 K c a l

2 N a (s) + 2 H C l (g) \to 2 N a C l (s) + H_{2} (g)

Δ H = - 152 K c a l

N a (s) + \frac{1}{2} C l_{2} (g) \to N a C l (s)

Δ H = ?

Δ H > Δ U

Δ H < Δ U

C (g r a p h i t e) + O_{2} (g) ⟶ C O_{2} (g)

P C l_{5} (g) ⟶ P C l_{3} (g) + C l_{2} (g)

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