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Reactivity 1. What drives chemical reactions? - Energy from fuels

Grade 11IBChemistry

Review the key concepts, formulae, and examples before starting your quiz.

🔑Concepts

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Chemical energy is a form of potential energy stored in the chemical bonds of substances. The total energy content of a system at constant pressure is defined as Enthalpy (HH).

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A reaction is Exothermic when energy is released to the surroundings. In these reactions, the enthalpy of the products is lower than the enthalpy of the reactants (ΔH<0\Delta H < 0).

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A reaction is Endothermic when energy is absorbed from the surroundings. In these reactions, the enthalpy of the products is higher than the enthalpy of the reactants (ΔH>0\Delta H > 0).

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Standard Enthalpy Change of Combustion (ΔHc⊖\Delta H_c^{\ominus}) is the enthalpy change when one mole of a substance is completely burned in excess oxygen under standard conditions (298K298 K, 100kPa100 kPa).

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Specific Energy is the energy released per unit mass of fuel (usually kJ g−1kJ\,g^{-1} or MJ kg−1MJ\,kg^{-1}). It is calculated as: Specific Energy=ΔHcMolar Mass\text{Specific Energy} = \frac{\Delta H_c}{\text{Molar Mass}}.

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Energy Density is the energy released per unit volume of fuel (usually kJ dm−3kJ\,dm^{-3}). It is calculated as: Energy Density=Specific Energy×Density\text{Energy Density} = \text{Specific Energy} \times \text{Density}.

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Calorimetry is the experimental technique used to measure the heat energy (qq) exchanged during a chemical reaction using the equation q=mcΔTq = mc\Delta T.

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Temperature is a measure of the average kinetic energy of particles, whereas Heat is the total energy transferred due to a temperature difference.

📐Formulae

q=mcΔTq = mc\Delta T

ΔH=−qn\Delta H = \frac{-q}{n}

Specific Energy=ΔHcM\text{Specific Energy} = \frac{\Delta H_c}{M}

Energy Density=Specific Energy×ρ\text{Energy Density} = \text{Specific Energy} \times \rho

ΔHrxn⊖=∑ΔHf⊖(products)−∑ΔHf⊖(reactants)\Delta H_{rxn}^{\ominus} = \sum \Delta H_{f}^{\ominus}(\text{products}) - \sum \Delta H_{f}^{\ominus}(\text{reactants})

💡Examples

Problem 1:

A student burns 0.46 g0.46\,g of ethanol (C2H5OHC_2H_5OH) to heat 150.0 g150.0\,g of water in a copper calorimeter. The temperature of the water increases from 20.0∘C20.0^{\circ}C to 35.0∘C35.0^{\circ}C. Calculate the experimental enthalpy of combustion of ethanol in kJ mol−1kJ\,mol^{-1}. (cwater=4.18 J g−1K−1c_{water} = 4.18\,J\,g^{-1}K^{-1}, Molar mass of ethanol = 46.08 g mol−146.08\,g\,mol^{-1})

Solution:

  1. Calculate the heat absorbed by water (qq): q=mcΔTq = mc\Delta T q=150.0 g×4.18 J g−1K−1×(35.0−20.0)Kq = 150.0\,g \times 4.18\,J\,g^{-1}K^{-1} \times (35.0 - 20.0)K q=150.0×4.18×15.0=9405 J=9.405 kJq = 150.0 \times 4.18 \times 15.0 = 9405\,J = 9.405\,kJ

  2. Calculate the moles of ethanol burned (nn): n=mM=0.46 g46.08 g mol−1≈0.010 moln = \frac{m}{M} = \frac{0.46\,g}{46.08\,g\,mol^{-1}} \approx 0.010\,mol

  3. Calculate Enthalpy Change (ΔH\Delta H): ΔH=−qn=−9.405 kJ0.010 mol=−940.5 kJ mol−1\Delta H = \frac{-q}{n} = \frac{-9.405\,kJ}{0.010\,mol} = -940.5\,kJ\,mol^{-1}

Explanation:

We first determine the energy transferred to the water. Since the reaction is combustion (exothermic), the enthalpy change is negative. We divide the heat by the number of moles of fuel consumed to find the molar enthalpy.

Problem 2:

Compare the specific energy of Hydrogen (H2H_2, ΔHc=−286 kJ mol−1\Delta H_c = -286\,kJ\,mol^{-1}) and Methane (CH4CH_4, ΔHc=−890 kJ mol−1\Delta H_c = -890\,kJ\,mol^{-1}).

Solution:

For Hydrogen (H2H_2): M=2.02 g mol−1M = 2.02\,g\,mol^{-1} Specific Energy=286 kJ mol−12.02 g mol−1≈141.6 kJ g−1\text{Specific Energy} = \frac{286\,kJ\,mol^{-1}}{2.02\,g\,mol^{-1}} \approx 141.6\,kJ\,g^{-1}

For Methane (CH4CH_4): M=16.05 g mol−1M = 16.05\,g\,mol^{-1} Specific Energy=890 kJ mol−116.05 g mol−1≈55.5 kJ g−1\text{Specific Energy} = \frac{890\,kJ\,mol^{-1}}{16.05\,g\,mol^{-1}} \approx 55.5\,kJ\,g^{-1}

Explanation:

Specific energy is the energy per unit mass. Even though methane has a higher molar enthalpy of combustion, hydrogen has a much higher specific energy because its molar mass is significantly lower, making it a very lightweight fuel source.