Atmospheric Pressure and Gas Laws

Atmospheric Pressure - Pressure's Ups & Downs

• Weight of air column; sea level standard: 760 mmHg (1 atm). Directly impacts inspired partial pressure of oxygen ($P_I O_2$).
• Altitude Effects:
  • Pressure ↓ exponentially with ↑ altitude.
  • 18,000 ft (5500m): Barometric Pressure (P_B) ≈ 380 mmHg (1/2 atm).
  • 30,000 ft: P_B ≈ 226 mmHg (0.3 atm).
  • Armstrong limit (63,000 ft): P_B = 47 mmHg (water boils at body temp).
• Diving Effects (Depth):
  • Pressure ↑ by 1 atm per 10m (33ft) descent in seawater.
  • At 10m depth: Total pressure ≈ 2 atm (1 atm ambient + 1 atm water).
• Measured by barometer.

⭐ Barometric pressure halves to ~380 mmHg at 18,000 ft, a critical threshold for hypoxia due to reduced oxygen availability.

, 18,000 ft (380 mmHg), and Armstrong limit (47 mmHg))

Boyle's & Dalton's Laws - Squeeze & Share Show

• Boyle's Law (The "Squeeze")

  • States: Constant temp: a gas's volume is inversely proportional to its pressure.
  • Formula: $P_1V_1 = P_2V_2$
  • 📌 Mnemonic: "Boyle's: Pressure & Volume Inverse."
  • Clinical Relevance:
    • Barotrauma: Pressure changes cause volume shifts in gas-filled body cavities.
      • Descent (↑P, ↓V): Middle ear/sinus squeeze.
      • Ascent (↓P, ↑V): Lung overexpansion, dental barotrauma.
    • Air Embolism: Gas bubbles expand with ↓ ambient pressure (rapid ascent).
    • Pneumothorax: Volume increases with altitude (↓ ambient pressure).

• Dalton's Law (The "Share")

  • States: Total pressure of a gas mixture is sum of its components' partial pressures.
  • Formula: $P_{total} = P_1 + P_2 + ... + P_n$
  • 📌 Mnemonic: "Dalton's Divides Pressure Pie."
  • Clinical Relevance:
    • Altitude Hypoxia: ↓ Total atm. pressure → ↓ inspired $PO_2$ (e.g., 5500m: $P_{atm} \approx \textbf{380 mmHg}$, $PIO_2 \approx \textbf{80 mmHg}$).
    • Diving:
      • Nitrogen Narcosis: ↑ $PN_2$ at depth impairs cognition.
      • Oxygen Toxicity: ↑ $PO_2$ at depth (e.g., >1.4 ATA $O_2$) → seizures.
    • Gas Exchange: Drives $O_2$/$CO_2$ diffusion via partial pressure gradients.
  • ⭐ > At sea level (Atm. pressure = 760 mmHg), inspired $PO_2$ from dry air (21% $O_2$) is $\approx \textbf{159.6 mmHg}$ ($760 \times 0.21$).

Henry's Law & Other Gas Laws - Soluble Gases & Temp Tales

• Henry's Law: Dissolved gas $\propto P_{gas}$ above liquid.
  • Formula: $C = k \cdot P_{gas}$
  • ↑$P_{gas}$ → ↑Gas dissolved (e.g., N₂ in tissues (diving)).
• Temperature Effects:
  • Gases: ↑Temp → ↑Volume (Charles' Law) or ↑Pressure (Gay-Lussac's Law).
  • Solubility: ↑Temp → ↓Gas solubility in liquids (Critical for DCS).
• Key Soluble Gases (Diving):
  • N₂: Moderately soluble; narcosis, DCS.
  • CO₂: Highly soluble.
  • He: Low solubility; Heliox for deep dives: ↓narcosis.

⭐ Nitrogen narcosis ("rapture of the deep") is primarily caused by the increased partial pressure of nitrogen dissolving in neuronal membranes, typically significant below 30 m (4 ATA).

High‑Yield Points - ⚡ Biggest Takeaways

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