States of matter: arrangement, motion, energies of particles, changes of state, gas laws
States of Matter
Solid
• Arrangement: Particles are tightly packed in a regular lattice. • Motion: Vibrate about fixed positions. • Energy: Low kinetic energy; most energy is stored as potential energy in the lattice. • Analogy: Think of a crowded dance floor where everyone is standing still but can wiggle a little. 🎶
Liquid
• Arrangement: Particles are close but not fixed; they slide past each other. • Motion: Particles move freely, giving liquids their fluidity. • Energy: Intermediate kinetic energy; particles have enough energy to overcome some lattice forces. • Analogy: Like a busy highway where cars can change lanes but stay close together. 🚗
Gas
• Arrangement: Particles are far apart and move independently. • Motion: Rapid, random motion in all directions. • Energy: High kinetic energy; particles rarely interact. • Analogy: Imagine a crowded playground where everyone runs around freely. 🏃♂️
Changes of State
Melting & Freezing
• Melting (solid → liquid): Energy added as heat of fusion $\,\Delta H_f^\circ\,$. • Freezing (liquid → solid): Energy released as $\,\Delta H_f^\circ\,$. • Exam tip: Remember that the temperature remains constant during the phase change.
Vaporisation & Condensation
• Boiling (liquid → gas): Energy added as heat of vaporisation $\,\Delta H_v^\circ\,$. • Condensation (gas → liquid): Energy released as $\,\Delta H_v^\circ\,$. • Analogy: Think of boiling water as a party where people (molecules) leave the room (liquid) to dance in the open air (gas). 🎉 • Exam tip: Boiling point depends on pressure; lower pressure → lower boiling point.
Sublimation & Deposition
• Sublimation (solid → gas): Direct transition without passing through liquid. • Deposition (gas → solid): Direct reverse transition. • Example: Dry ice (solid CO₂) sublimates at room temperature. ❄️ • Exam tip: Use phase diagrams to determine conditions for sublimation.
Gas Laws
| Law | Equation | Key Points |
|---|---|---|
| Boyle’s Law | $P_1V_1 = P_2V_2$ | Pressure inversely proportional to volume at constant temperature. |
| Charles’ Law | $V_1/T_1 = V_2/T_2$ | Volume directly proportional to temperature (Kelvin) at constant pressure. |
| Avogadro’s Law | $V_1/n_1 = V_2/n_2$ | Equal volumes contain equal numbers of molecules at same conditions. |
| Ideal Gas Law | $PV = nRT$ | Combines all three laws; $R = 0.0821\,\text{L·atm·mol}^{-1}\text{K}^{-1}$. |
Exam Tip: Always check units before substituting values into the ideal gas equation. Convert pressures to atm, volumes to L, and temperatures to K.
🧪 Remember that real gases deviate from ideal behaviour at high pressures and low temperatures; use the Van der Waals equation if required.
Practice Question
A sample of 2.00 mol of an ideal gas occupies 22.4 L at 25 °C and 1.00 atm.
- Calculate the pressure if the volume is reduced to 11.2 L while keeping temperature constant.
- What will be the new volume if the pressure is increased to 2.00 atm at the same temperature?
Quick Flashcards
- Solid: Fixed shape, fixed volume, low kinetic energy.
- Liquid: Fixed volume, takes shape of container, moderate kinetic energy.
- Gas: Variable shape, variable volume, high kinetic energy.
- Phase change: Temperature stays constant; energy is absorbed or released as latent heat.
- Ideal Gas Law: $PV = nRT$ – remember the constant $R$ and unit conversions.
Revision
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