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Structure 1.5 Flashcards

Active Recall for Gas Laws

1.5.1 Gas Laws: read all 6 questions and answers
  1. State Boyle's Law.
    At constant temperature, the pressure of a fixed mass of gas is inversely proportional to its volume. P1V1 = P2V2.
  2. State Charles's Law.
    At constant pressure, the volume of a fixed mass of gas is directly proportional to its absolute temperature (in Kelvin). V1/T1 = V2/T2.
  3. State Gay-Lussac's Law.
    At constant volume, the pressure of a fixed mass of gas is directly proportional to its absolute temperature (in Kelvin). P1/T1 = P2/T2.
  4. State Avogadro's Law.
    Equal volumes of gases at the same temperature and pressure contain equal numbers of particles (moles). V is proportional to n.
  5. What is STP in the IB?
    Standard Temperature and Pressure: 273.15 K (0 degrees C) and 100 kPa.
  6. What is the combined gas law?
    P1V1/T1 = P2V2/T2 (combines Boyle's, Charles's, and Gay-Lussac's laws for a fixed amount of gas).
1.5.2 Ideal Gas Equation: read all 6 questions and answers
  1. State the ideal gas equation.
    PV = nRT, where P = pressure (Pa), V = volume (m³), n = moles, \(R = 8.314\text{ J mol}^{-1}\text{ K}^{-1}\), T = temperature (K).
  2. What are the SI units for PV = nRT?
    P in Pa (not kPa), V in m³ (not dm³ or cm³), T in K (not degrees C), \(R = 8.314\text{ J mol}^{-1}\text{ K}^{-1}\).
  3. How do you convert kPa to Pa?
    Multiply by 1000. So 101.3 kPa = 101300 Pa.
  4. How do you convert dm³ to m³?
    Divide by 1000. So 22.7 dm³ = 0.0227 m³.
  5. How can you use PV = nRT to find molar mass?
    Since n = m/M, rearrange to M = mRT/(PV). Measure mass, temperature, pressure, and volume to find M.
  6. What is the value of the gas constant R?
    \(R = 8.314\text{ J mol}^{-1}\text{ K}^{-1}\). This is given in the IB data booklet.
1.5.3 Real vs Ideal Gases: read all 6 questions and answers
  1. What are the two key assumptions of the ideal gas model?
    1. Gas particles have negligible volume compared to the container. 2. There are no intermolecular forces between particles.
  2. Under what conditions do real gases behave most ideally?
    At high temperature (particles have high kinetic energy to overcome IMFs) and low pressure (particles are far apart, so volume and IMFs are negligible).
  3. Why do real gases deviate from ideal behaviour at high pressure?
    At high pressure, particles are forced close together. Their actual volume becomes significant compared to container volume, and IMFs become stronger.
  4. Why do real gases deviate from ideal behaviour at low temperature?
    At low temperature, particles move slowly and have less kinetic energy. IMFs become significant relative to kinetic energy, reducing the effective pressure.
  5. Which type of gas behaves most ideally?
    Small, non-polar gases (e.g. He, H2) because they have the weakest IMFs and smallest volume. Large polar molecules deviate the most.
  6. What happens to all real gases at sufficiently low temperature?
    They liquefy (condense), which an ideal gas would never do. This is because IMFs become dominant over kinetic energy.
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