Consider the following information. Ethanol burns in excess air according to the following equation. C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(g) ΔH = −1364 kJ mol−1 The cooking pot is made from aluminium and has a mass of 150 g. The specific heat capacity of aluminium is 9.000 J g−1 °C−1. The specific heat capacity of water is 4.181 J g−1 °C−1. i. Calculate the minimum amount of energy, in kJ, required to heat 550 g of water and the pot from 18.5°C to 100.0°C. ii. Calculate the mass, in g, of ethanol that needs to be completely burnt to provide this energy. iii. Only 35% of the energy released by the combustion of ethanol is transferred to the cooking pot and contents. Calculate the mass, in g, of ethanol that needs to be burnt in practice to heat the water and the pot from 18.5°C to 100.0°C. b. Other camping stoves use butane (C4H10) as fuel. Given that, on complete combustion, 6.00 g of butane releases the same amount of energy as 10.0 g of ethanol, calculate the magnitude of ΔH, in kJ mol−1, for the reaction 2C4H10(g) + 13O2(g) → 8CO2(g) + 10H2O(g)

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Consider the following information - VCE - SSCE Chemistry - Question 6 - 2007 - Paper 1

Question 6

Consider the following information. Ethanol burns in excess air according to the following equation. C2H5OH(l) + 3O2(g) → 2CO2(g) + 3H2O(g) ΔH = −1364 kJ mol−1 Th... show full transcript

Worked Solution & Example Answer:Consider the following information - VCE - SSCE Chemistry - Question 6 - 2007 - Paper 1

Step 1

Calculate the minimum amount of energy, in kJ, required to heat 550 g of water and the pot from 18.5°C to 100.0°C.

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Answer

To calculate the energy required, we can use the formula:

$Q = mc\Delta T$

where:

$m$ is the mass (g)
$c$ is the specific heat capacity (J g−1 °C−1)
$\Delta T$ is the change in temperature (°C)

Calculate the energy needed to heat the water:
- Mass of water, $m_w = 550 \, \text{g}$
- Specific heat capacity, $c_w = 4.181 \, \text{J g}^{-1} °C^{-1}$
- Change in temperature, $\Delta T_w = 100.0 - 18.5 = 81.5 \, °C$
- Energy, $Q_w = m_w c_w \Delta T_w = 550 imes 4.181 imes 81.5 = 198,242.25 \, \text{J}$
Calculate the energy needed to heat the pot:
- Mass of pot, $m_p = 150 \, \text{g}$
- Specific heat capacity, $c_p = 9.000 \, \text{J g}^{-1} °C^{-1}$
- Change in temperature, $\Delta T_p = 100.0 - 18.5 = 81.5 \, °C$
- Energy, $Q_p = m_p c_p \Delta T_p = 150 imes 9.000 imes 81.5 = 1,225,500 \, \text{J}$
Total energy:
- $Q_{total} = Q_w + Q_p = 198,242.25 + 1,225,500 = 1,423,742.25 \, \text{J}$
Convert to kJ:
- $\frac{1,423,742.25}{1000} = 1423.74 \, \text{kJ}$

Step 2

Calculate the mass, in g, of ethanol that needs to be completely burnt to provide this energy.

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Answer

Using the enthalpy change for ethanol combustion:

$\Delta H = -1364 \, \text{kJ mol}^{-1}$

Energy needed is approximately 1423.74 kJ.
Calculate moles of ethanol required:
- Moles of ethanol, $n = \frac{Q}{|\Delta H|} = \frac{1423.74}{1364} = 1.04 \, \text{mol}$
Calculate mass of ethanol:
- Molar mass of ethanol (C2H5OH) = 46.07 g/mol
- Mass = $n \times \text{Molar Mass} = 1.04 \times 46.07 = 47.94 \, \text{g}$ .

Step 3

Only 35% of the energy released by the combustion of ethanol is transferred to the cooking pot and contents. Calculate the mass, in g, of ethanol that needs to be burnt in practice to heat the water and the pot from 18.5°C to 100.0°C.

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Answer

Calculate the effective energy used:
- Effective energy = $0.35 \times 1423.74 = 497.31 \, \text{kJ}$ .
Calculate the moles of ethanol needed for this energy:
- $n = \frac{497.31}{1364} = 0.365 \, \text{mol}$ .
Calculate the mass of ethanol:
- Mass = $0.365 \times 46.07 = 16.80 \, \text{g}$ .

Step 4

Calculate the magnitude of ΔH, in kJ mol−1, for the reaction.

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Answer

Using the fact that 10.0 g of ethanol releases the same energy as 6.00 g of butane:

Calculate energy released by 10.0 g of ethanol:
- $\Delta H = 1364 \, \text{kJ mol}^{-1}$ is the energy for 1 mol.
- For 10.0 g: Moles of ethanol = $\frac{10}{46.07} = 0.217 \, \text{mol}$ , then energy = $0.217 × 1364 = 296.9 \, \text{kJ}$ .
Since 6.00 g of butane releases the same energy, we can set up:
- $\Delta H$ for butane corresponds to this energy, confirmed through stoichiometry in the reaction.

Consider the following information - VCE - SSCE Chemistry - Question 6 - 2007 - Paper 1

Worked Solution & Example Answer:Consider the following information - VCE - SSCE Chemistry - Question 6 - 2007 - Paper 1

Calculate the minimum amount of energy, in kJ, required to heat 550 g of water and the pot from 18.5°C to 100.0°C.

Calculate the mass, in g, of ethanol that needs to be completely burnt to provide this energy.

Only 35% of the energy released by the combustion of ethanol is transferred to the cooking pot and contents. Calculate the mass, in g, of ethanol that needs to be burnt in practice to heat the water and the pot from 18.5°C to 100.0°C.

Calculate the magnitude of ΔH, in kJ mol−1, for the reaction.

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