3.3 States of Matter, Thermal Energy & Phase Changes

3.3 States of Matter, Thermal Energy & Phase Changes

GED Exam Core Concept: Matter exists in distinct states defined by the energy and arrangement of its constituent particles. On the GED Science exam, you will analyze phase diagrams and heating curves, differentiate between heat and temperature, and evaluate endothermic versus exothermic energy transfers.

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The Kinetic Molecular Theory (KMT)

The Kinetic Molecular Theory (KMT) provides a microscopic model explaining the physical behavior of matter:

1. Continuous Motion: All matter is made of tiny particles (atoms or molecules) that are in constant, random motion.
2. Kinetic Energy & Temperature: The average kinetic energy ($KE = \frac{1}{2}mv^2$) of particles is directly proportional to the absolute temperature of the substance in Kelvin ($K$). Higher temperature = faster moving particles.
3. Intermolecular Forces: Particles exert attractive forces on one another. The balance between particle kinetic energy and intermolecular attractions determines whether a substance is a solid, liquid, or gas.

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Characteristics of the Four States of Matter

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Thermal Energy vs. Temperature vs. Heat

It is essential on the GED exam to distinguish between these three related thermal physics terms:

1. Temperature ($T$):

   • A measure of the average kinetic energy of the individual particles in a sample.
   • Measured using a thermometer in degrees Celsius ($^\circ\text{C}$), Fahrenheit ($^\circ\text{F}$), or Kelvin ($\text{K}$).
   • Independent of sample size (an intensive property). A cup of boiling water has the same temperature as a bathtub of boiling water ($100^\circ\text{C}$).

2. Thermal Energy:

   • The total energy (kinetic energy of motion + potential energy of intermolecular bonds) of all particles in a sample.
   • Dependent on sample size (an extensive property). A bathtub of boiling water contains much more thermal energy than a cup of boiling water.

3. Heat ($Q$):

   • Thermal energy in transit flowing spontaneously from a body of higher temperature to a body of lower temperature.
   • Measured in Joules ($\text{J}$) or calories ($\text{cal}$).

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Phase Changes & Phase Transitions

A phase change is a physical transition of a substance from one state of matter to another, caused by adding or removing thermal energy.

Six Fundamental Phase Transitions:

1. Melting (Fusion): Solid $\rightarrow$ Liquid (Absorbs heat - Endothermic)
2. Freezing (Solidification): Liquid $ ightarrow$ Solid (Releases heat - Exothermic)
3. Vaporization (Boiling/Evaporation): Liquid $ ightarrow$ Gas (Absorbs heat - Endothermic)
4. Condensation: Gas $ ightarrow$ Liquid (Releases heat - Exothermic)
5. Sublimation: Solid $ ightarrow$ Gas directly (Absorbs heat - Endothermic, e.g., Dry Ice $\text{CO}_2$)
6. Deposition: Gas $ ightarrow$ Solid directly (Releases heat - Exothermic, e.g., Frost on window)

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Endothermic vs. Exothermic Phase Changes

• Endothermic Phase Changes: Processes that absorb energy from the surrounding environment to break or weaken intermolecular bonds ($Q > 0$). $\text{Solid} \xrightarrow[\text{Melting}]{+\text{Heat}} \text{Liquid} \xrightarrow[\text{Vaporization}]{+\text{Heat}} \text{Gas}$

• Exothermic Phase Changes: Processes that release energy into the surrounding environment as intermolecular bonds form ($Q < 0$). $\text{Gas} \xrightarrow[\text{Condensation}]{-\text{Heat}} \text{Liquid} \xrightarrow[\text{Freezing}]{-\text{Heat}} \text{Solid}$

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Heating Curves: Why Temperature Plateaus During Phase Changes

When a solid substance is heated at a constant rate, its temperature changes as shown on a standard heating curve:

Temperature (°C)
     |
 100 |                       / (Gas Heating)
     |                      / 
 100 |..........+----------+ (Boiling / Vaporization Plateau: Liquid + Gas)
     |         / 
  0  |        / (Liquid Heating)
  0  |...+---+ (Melting Plateau: Solid + Liquid)
     |  /
     | / (Solid Heating)
  ---+--------------------------------------------> Heat Added / Time

Critical GED Graph Interpretation Rule:

• Sloped Segments: Added heat increases the kinetic energy (speed) of particles $\rightarrow$ Temperature increases ($Q = m c \Delta T$).
• Horizontal Flat Plateaus: Added heat does NOT increase particle speed or temperature. Instead, temperature remains strictly constant while heat energy is consumed as potential energy to break intermolecular attractions and change the phase ($Q = m L$).
  • Melting Point Plateau: Solid and liquid co-exist in equilibrium ($0^\circ\text{C}$ for pure water).
  • Boiling Point Plateau: Liquid and gas co-exist in equilibrium ($100^\circ\text{C}$ for pure water).

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Quantitative Thermal Calculations

1. Temperature Change Equation (No Phase Transition):

$Q = m \cdot c \cdot \Delta T$ Where:

• $Q$ = Heat absorbed or released (Joules, $\text{J}$)
• $m$ = Mass of substance (grams, $\text{g}$)
• $c$ = Specific heat capacity ($\text{J}/\text{g}^\circ\text{C}$)
• $\Delta T$ = Change in temperature ($T_{\text{final}} - T_{\text{initial}}$)

2. Phase Change Latent Heat Equation (At Constant Temperature):

$Q = m \cdot L$ Where:

• $L_f$ = Latent Heat of Fusion (for melting/freezing; $334\text{ J/g}$ for water)
• $L_v$ = Latent Heat of Vaporization (for boiling/condensation; $2260\text{ J/g}$ for water)

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Worked Examples

Worked Example: Heat Calculation for Phase Transition

Problem: How much thermal energy in Joules is required to melt a $50.0\text{ g}$ block of ice completely into liquid water at its melting point of $0^\circ\text{C}$? (Latent heat of fusion of water $L_f = 334\text{ J/g}$).

Solution:

1. Select Equation: Because this is a phase change at constant temperature ($0^\circ\text{C}$), use $Q = m \cdot L_f$.
2. Substitute Values: $Q = 50.0\text{ g} \times 334\text{ J/g} = \mathbf{16,700\text{ Joules}}$