Question

How much heat is released when 6.00 g of methane is combusted using the reaction below? CH4 + 2O2 ? CO2 + 2H2O ?Hrxn = -890. KJ

Answer 1

Answer : The amount of heat released is, -121.04 KJ

Solution : Given,

Enthalpy of reaction, $\Delta H_{rxn}$ = -890 KJ

Mass of methane = 6 g

Molar mass of methane = 44 g/mole

First we have to calculate the moles of methane.

$\text{Moles of methane}=\frac{\text{Mass of methane}}{\text{Molar mass of methane}}=\frac{6g}{44g/mole}=0.136moles$

The given balanced reaction is,

$CH_4+2O_2\rightarrow CO_2+2H_2O$

From the balanced reaction we conclude that

1 mole of methane releases heat = -890 KJ

0.136 moles of methane releases heat = $\frac{0.136moles}{1mole}\times -890KJ=-121.04KJ$

Therefore, the amount of heat released is, -121.04 KJ

Answer 2

Answer:

-333.75 KJ

Explanation:

From the reaction CH₄ + 2O₂ ⇒ CO₂ + 2H₂O    ΔHrxn = -890 KJ.

From the reaction, 1 mole of CH₄ combusts to give -890 KJ.

We now calculate the number of moles, n of CH₄ in 6.00g

n = mass of CH₄/molar mass of CH₄

molar mass of CH₄ = 12 g/mol + 4 × 1 g/mol = (12 + 4) g/mol = 16 g/mol

mass of CH₄ = 6.00 g

n = 6.00 g /16 g/mol = 0.375 mol

So we have 0.375 mol of CH₄ in 6.00g

From the reaction, 1 mole of CH₄ produces -890 KJ

then 0.375 mol produces 0.375 mol × -890 KJ/1 mol = -333.75 KJ