When butane (C₄H₁₀, 58.14 g/mol) is burned in oxygen according to the reaction below, 5755 kJ of heat is released.About how much butane should a backpacker carry if s/he needs to heat 1.50 L of water from 17$\degree$C to 100$\degree$C? Assume all the energy from the combustion is used to raise the water's temperature.[c_pwater) = 4.18 J/(g . $\degree$C); d(water) = 1.00g/mL; c_p(butane) =1.69 J/(g.$\degree$C)] 2 C₄ H₁₀ (g) + 13O₂ (g) $\rightarrow$ 8 CO₂ (g) + 10 H₂ O(l)

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Quizplus · Accepted Answer

First, calculate the energy needed to heat 1.50 L of water from 17°C to 100°C. The mass of the water is 1500 g (since 1 mL = 1 g and there are 1000 mL in 1 L). The temperature change (ΔT) is 100°C - 17°C = 83°C. The specific heat capacity of water (c) is 4.18 J/g°C. The energy (q) required is calculated using the formula q = mcΔT, which gives q = 1500 g * 4.18 J/g°C * 83°C = 519,090 J or 519.09 kJ.Given that the combustion of butane releases 5755 kJ, we can find the amount of butane needed to produce 519.09 kJ by setting up a proportion: 5755 kJ is to the molar mass of butane (58.14 g/mol) as 519.09 kJ is to x grams of butane. Solving for x gives x = (519.09 kJ / 5755 kJ) * 58.14 g/mol ≈ 11 g.

When Butane (C4H10, 58 °\degree° C to 100 °\degree° C? Assume All the Energy from the Combustion Is Used

When Butane (C₄H₁₀, 58 $\degree$ C to 100 $\degree$ C? Assume All the Energy from the Combustion Is Used