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Equilibrium Constant and Standard Free Energy Change

Δf H⊖ per mol...

Question

$Δ_{f} H^{⊖}$ per mole of $N H_{3} (g)$ , $N O (g)$ and $H_{2} O (l)$ are $- 11.04$ , $+ 21.60$ and $- 68.32 K c a l$ , respectively. Calculate the standard heat of reaction at constant pressure and at a constant volume for the reaction:
$4 N H_{3} (g) + 5 O_{2} (g) ⟶ 4 N O (g) + 6 H_{2} O (l)$

Δ H = - 279.36 K c a l

Δ U = - 276.38 K c a l

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Δ H = + 279.36 K c a l

Δ U = - 276.38 K c a l

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Δ H = - 279.36 K c a l

Δ U = + 276.38 K c a l

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None of these

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Solution

The correct option is B $Δ H = - 279.36 K c a l$ $Δ U = - 276.38 K c a l$ .

$Δ n_{g} = 4 - 9$ $= - 5$
$Δ_{f} H^{⊖} O_{2} = 0$ (definition)
$Δ H = [4 Δ_{f} H^{⊖} (N O) + 6 Δ_{f} H^{⊖} (H_{2} O)] - 4 Δ_{f} H^{⊖} (N H_{3})$ $= 4 \times 21.6 + 6 \times - (68.32 - (- 11.04)$ $= - 279.36 K c a l$ .
$Δ U = Δ H - Δ n_{g} R T$ $= - 279.36 - (- 5) \times 2 \times 10^{- 3} \times 298$ $= - 276.38 K c a l$ .

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Similar questions

For the conversion of 1 mole of $S O_{2} (g)$ into $S O_{3} (g)$ the enthalpy of reaction at constant volume, $Δ U$ at $298 K$ is $- 97.027 K J$ . What is the enthalpy of reaction, $Δ H$ at constant pressure?

Q. Heat of reaction at constant pressure and heat of reaction at constant volume for the gaseous reaction

N_{2} + 3 H_{2} ⟶ 2 N H_{3}

differ

(Δ H - Δ U)

by the amount:

Q. For the reaction

C_{3} H_{8} (g) + 5 O_{2} (g) \to 3 {C O}_{2} (g) + 4 H_{2} O (l)

at constant temperature,

Δ H - Δ U

is _______ .

Q. The following sequence of reaction occurs in commercial production of aqueous nitric acid.

4 N H_{3} (g) + 5 O_{2} (g) \to 4 N O (g) + 6 H_{2} O (l) Δ H = - 904 k J . . . (1)

2 N O (g) + O_{2} (g) \to 2 N O_{2} (g) Δ H = - 112 k J . . . (2)

3 N O_{2} (g) + H_{2} O (l) \to 2 H N O_{3} (a q) + N O (g) Δ H = - 140 k J . . . (3)

Determine the magnitude of total heat

(i n kJ/mole)

liberated at constant pressure for the production of exactly 1 mole of aqueous nitric acid by this process

Q. For the conversion of

1

mole of

{S O}_{2} (g)

into

{S O}_{3} (g)

the enthalpy of reaction at constant volume,

Δ U

, at

298 K

- 97.027 k J

. Calculate the enthalpy of reaction,

Δ H

, at constant pressure.

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ΔfH⊖ per mole of NH3(g), NO(g) and H2O(l) are −11.04, +21.60 and −68.32Kcal, respectively. Calculate the standard heat of reaction at constant pressure and at a constant volume for the reaction:4NH3(g)+5O2(g)⟶4NO(g)+6H2O(l)

The correct option is B ΔH=−279.36Kcal ΔU=−276.38Kcal. Δng=4−9 =−5 ΔfH⊖O2=0 (definition)ΔH=[4ΔfH⊖(NO)+6ΔfH⊖(H2O)]−4ΔfH⊖(NH3) =4×21.6+6×−(68.32−(−11.04) =−279.36Kcal. ΔU=ΔH−ΔngRT =−279.36−(−5)×2×10−3×298 =−276.38Kcal.