Examples of Phase Changes and Enthalpy

Understanding Phase Changes and Enthalpy

Phase changes refer to the transformation of matter from one state to another, such as solid to liquid or liquid to gas. Enthalpy, a measure of total energy in a thermodynamic system, plays a crucial role in these changes. This article presents three practical examples of phase changes and enthalpy to illustrate these concepts effectively.

Example 1: Melting Ice - Fusion Process

In everyday life, melting ice is a common occurrence, especially during warm weather. When ice at 0°C absorbs heat, it undergoes a phase change from solid to liquid, which is known as fusion. The heat absorbed during this process is quantified as the enthalpy of fusion. This example is fundamental in understanding how energy is transferred during phase changes.

For instance, consider a 100 g block of ice at 0°C. The enthalpy of fusion for ice is approximately 334 J/g. Thus, to calculate the total heat energy required to melt the ice completely, the formula is:

Q = m * ΔH_f
where:

• Q = heat energy absorbed (in joules)
• m = mass of the ice (in grams)
• ΔH_f = enthalpy of fusion (in J/g)

Calculating this gives:

Q = 100 g * 334 J/g = 33,400 J

This means that to melt 100 g of ice at 0°C into water at the same temperature, 33,400 joules of energy must be absorbed.

Notes:

• The temperature remains constant during the melting process until the entire ice block is converted to water.
• If you were to cool the water back down to 0°C, the same amount of energy would be released, demonstrating conservation of energy.

Example 2: Boiling Water - Vaporization Process

Boiling water is a well-known example of a phase change from liquid to gas, referred to as vaporization. When water reaches its boiling point at 100°C under standard atmospheric pressure, it begins to convert into steam. The enthalpy of vaporization is the energy required for this transformation.

For example, consider 250 g of water boiling at 100°C. The enthalpy of vaporization for water is about 2260 J/g. To find the total energy needed to vaporize this mass of water, we use the same formula:

Q = m * ΔH_v
where:

• ΔH_v = enthalpy of vaporization (in J/g)

Calculating this yields:

Q = 250 g * 2260 J/g = 565,000 J

Thus, 565,000 joules of energy are required to convert 250 g of water at 100°C into steam at the same temperature.

Notes:

• Like melting, the boiling process occurs at a constant temperature until all the water is converted to steam.
• The enthalpy of vaporization is significantly higher than that of fusion, reflecting the energy needed to overcome intermolecular forces in the liquid state.

Example 3: Condensation of Steam - Heat Release

The condensation of steam back into water is a crucial process in many industrial applications, such as in power plants. When steam cools down and transitions from a gas back to a liquid, it releases energy, quantified as the enthalpy of condensation, which is equal in magnitude but opposite in sign to the enthalpy of vaporization.

If 200 g of steam at 100°C condenses, the enthalpy of condensation for water is approximately -2260 J/g (the negative sign indicates energy release). To calculate the energy released during this process, we can again use our formula:

Q = m * ΔH_c
where:

• ΔH_c = enthalpy of condensation (in J/g)

This gives us:

Q = 200 g * (-2260 J/g) = -452,000 J

This indicates that 452,000 joules of energy are released when 200 g of steam condenses into water at 100°C.

Notes:

• The energy released during condensation can be harnessed for heating purposes, making this process particularly useful in energy recovery systems.
• The temperature remains constant during condensation until all steam has transformed into water.

In summary, the examples provided illustrate the critical relationship between phase changes and enthalpy in various contexts, showcasing how energy is absorbed or released during these transformations.