Butane (c4h10) undergoes combustion in excess oxygen to generate gaseous carbon dioxide and water. given δh°f[c4h10 (g) ] = - 124.7 kj/mol, δh°f[co2 (g) ] = - 393.5 kj/mol, δh°f[h2o (g) ] = - 241.8 kj/mol, how much energy is released (kj) when 8.30 g of butane is burned?

The value of Δ H butane (g) = - 124.7 kJ/mol

The value of Δ H CO2 (g) = - 393.5 kJ/mol

The value of Δ H H2O (g) = - 241.8 kJ/mol

Mass of butane, m = 8.30 gm

Molar mass of butane is 58 gm/mol

Consider the reaction,

C₄H₁₀ + 6.5 O₂ = 4CO₂ + 5H₂O

Calculating the value of Δ H° rxn:

ΔH°rxn = ∑nH° f (products) - ∑nH° f (reactants)

Substituting the values we get,

Δ H° rxn = 4 (-393.5) + 5 (-241.8) - (-124.7)

= - 1574 - 1209 + 124.7

= - 2783 - 124.7

= - 2658.3 kJ/mol

Now, calculate the number of moles of butane in 8.30 gm.

Number of moles = mass/molar mass

= 8.30 / 58

= 0.143 moles

Thus, the total energy released in the reaction is,

Q = number of moles * ΔH° rxn

= 0.143 * (2658.3)

= 380.14 kJ

Hence, the total heat released in the reaction is 380.14 kJ.