Calculating Enthalpy (part 2) energy is consumed by thawing 4.3 g ice? Use the chart below 1. H Constants for water 40.65 kJ/mol 285 83 kJ/mol 6 03 kJ/mol specific heat 4 186 J/g c molar mass 18 02 g 2. How much energy is generated from freezing 2.5 g water? Use the chart above. 3 How much energy is released when 6.0 g of water is condensed from water vapor?
The Correct Answer and Explanation is :
To determine the energy changes associated with phase transitions of water, we can use the provided constants and the relationships between mass, molar mass, and enthalpy changes. Here’s a step-by-step explanation for each scenario:
1. Energy Consumed by Thawing 4.3 g of Ice
When ice melts, it absorbs energy equal to its enthalpy of fusion. The enthalpy of fusion for water is 6.03 kJ/mol. To find the energy required to melt 4.3 g of ice:
- Calculate the number of moles of ice: [
\text{Moles of ice} = \frac{\text{mass}}{\text{molar mass}} = \frac{4.3\, \text{g}}{18.02\, \text{g/mol}} \approx 0.2386\, \text{mol}
] - Calculate the energy absorbed: [
\text{Energy} = \text{moles} \times \Delta H_{\text{fusion}} = 0.2386\, \text{mol} \times 6.03\, \text{kJ/mol} \approx 1.44\, \text{kJ}
]
Therefore, approximately 1.44 kJ of energy is consumed to thaw 4.3 g of ice.
2. Energy Released from Freezing 2.5 g of Water
Freezing is the reverse process of melting, and it releases energy equal to the enthalpy of fusion. To find the energy released when 2.5 g of water freezes:
- Calculate the number of moles of water: [
\text{Moles of water} = \frac{2.5\, \text{g}}{18.02\, \text{g/mol}} \approx 0.1387\, \text{mol}
] - Calculate the energy released: [
\text{Energy} = \text{moles} \times \Delta H_{\text{fusion}} = 0.1387\, \text{mol} \times 6.03\, \text{kJ/mol} \approx 0.84\, \text{kJ}
]
Thus, approximately 0.84 kJ of energy is released when 2.5 g of water freezes.
3. Energy Released when 6.0 g of Water Vapor Condenses
Condensation is the process where water vapor turns into liquid water, releasing energy equal to the enthalpy of vaporization. The enthalpy of vaporization for water is 40.65 kJ/mol. To find the energy released when 6.0 g of water vapor condenses:
- Calculate the number of moles of water vapor: [
\text{Moles of water vapor} = \frac{6.0\, \text{g}}{18.02\, \text{g/mol}} \approx 0.3330\, \text{mol}
] - Calculate the energy released: [
\text{Energy} = \text{moles} \times \Delta H_{\text{vaporization}} = 0.3330\, \text{mol} \times 40.65\, \text{kJ/mol} \approx 13.54\, \text{kJ}
]
Therefore, approximately 13.54 kJ of energy is released when 6.0 g of water vapor condenses.
Explanation:
The energy changes during phase transitions are due to the breaking or forming of intermolecular forces. During melting (fusion), energy is absorbed to overcome the hydrogen bonds holding water molecules in a solid structure, allowing them to move freely as a liquid. Conversely, during freezing, these bonds reform, releasing the same amount of energy.
Similarly, during vaporization, a significant amount of energy is required to completely separate water molecules from the liquid phase into the gas phase, overcoming intermolecular attractions. This is why the enthalpy of vaporization is substantially higher than the enthalpy of fusion. When condensation occurs, the process reverses, and the energy is released as water molecules transition from gas to liquid, reforming intermolecular attractions.
These enthalpy values are intrinsic properties of water and are crucial in understanding energy exchanges in various natural and industrial processes, such as climate dynamics, refrigeration, and heating systems.