MCAT topics → Fluids in Circulation and Gas Exchange → Partial Pressures (Dalton's Law)

Partial Pressures (Dalton's Law)

MCAT trap: Equates partial pressure with gas concentration rather than mole fraction times total pressure. Partial pressure equals the mole fraction of the gas multiplied by total pressure (P_i = χ_i × P_total), not its concentration directly.

Partial pressures show up constantly in MCAT passages on respiratory physiology, gas exchange, and altitude physiology, and they carry a high-yield misconception: at altitude, the fraction of oxygen in air is still ~21% — what drops is total atmospheric pressure, so PO₂ drops proportionally. Students who think thin air means less oxygen as a percentage get this backwards and miss clinical reasoning questions about hypoxia. The core mechanics: each gas's partial pressure is its mole fraction times the total pressure, and Dalton's law says the total is just their sum.

The MCAT hits this concept from three angles. First, pure recall and definition — can you state Dalton's law and write the formula? Second, calculation — given mole fractions or percentages and a total pressure, can you compute the partial pressure of a specific gas? Third, and most commonly in passages, cross-disciplinary application — connecting partial pressures to alveolar gas exchange, oxygen loading at the lungs, and how altitude changes PO₂ even though the air composition stays the same. That third angle is where students lose points because it requires you to integrate gas law mechanics with physiology.

The misconceptions here are sneaky. Students often conflate partial pressure with concentration, or assume that gas molecules in a mixture somehow interact and reduce total pressure. The altitude one is especially high-yield: most students wrongly think thin air means less oxygen as a fraction of air, when actually the O₂ percentage is still ~21% — it's the drop in total atmospheric pressure that reduces PO₂ and causes hypoxia. Get that distinction sharp before test day.

Common misconceptions

Common mistake
Wrong: The partial pressure of a gas equals its concentration in the mixture.
Right: Partial pressure equals the mole fraction of the gas multiplied by total pressure (P_i = χ_i × P_total), not its concentration directly.
Partial pressure and concentration are related but not equal. Partial pressure is calculated as mole fraction × total pressure (P_i = χ_i × P_total) — it depends on how many moles of that gas are present relative to total moles AND on the total pressure of the system. Concentration (like mol/L) doesn't account for total pressure. Two gas samples could have the same oxygen concentration but different PO₂ values if they're at different total pressures — this is exactly why altitude matters physiologically.
Common mistake
Wrong: At high altitude, the fraction of oxygen in air decreases, causing hypoxia.
Right: The fraction of O₂ in air remains ~21% at altitude, but total pressure decreases, so the partial pressure of O₂ (PO₂) falls, causing hypoxia.
Air at sea level and air at 15,000 feet have essentially the same composition — about 21% O₂ either way. What changes with altitude is total atmospheric pressure, which drops because there's less air sitting above you. Since PO₂ = 0.21 × P_total, when P_total drops, PO₂ drops proportionally. Less PO₂ means less driving force for O₂ to diffuse across the alveolar membrane into blood, causing hypoxia. The gas fraction stays fixed; the total pressure — and therefore each partial pressure — is what falls.
Common mistake
Wrong: Gases in a mixture interact so that the total pressure is less than the sum of individual partial pressures.
Right: For ideal gases, Dalton's law states that total pressure equals the sum of all partial pressures, with no interaction between gas molecules assumed.
Dalton's law applies to ideal gases, which by definition do not interact with each other. Each gas molecule moves independently, so each contributes its own pressure to the container walls as if the other gases weren't there. The total pressure is simply the sum of these independent contributions — there's no subtraction for 'interactions.' If you're seeing a scenario where total pressure seems less than expected, it's more likely a real-gas deviation or a phase change (like water vapor condensing), not gas-gas interaction reducing pressure.
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What the exam tests

  1. Know the definition of Dalton's law: the total pressure of a gas mixture equals the sum of the partial pressures of each individual gas (P_total = P₁ + P₂ + P₃ + ...).
  2. Calculate the partial pressure of a specific gas given its mole fraction (or percentage by volume) and the total pressure using P_i = χ_i × P_total.
  3. Apply partial pressure concepts to biological gas exchange — explain how alveolar PO₂ is determined, how O₂ loading at the lungs depends on PO₂, and why high altitude reduces PO₂ and causes hypoxia even though the fraction of O₂ in air remains constant at ~21%.

Can you avoid these mistakes?

A gas mixture contains 0.4 mol N₂, 0.1 mol O₂, and 0.5 mol CO₂ in a container at 800 mmHg total pressure. What is the partial pressure of O₂? Walk through your calculation using mole fractions.
At sea level (P_total ≈ 760 mmHg), PO₂ in alveolar air is approximately 100 mmHg. A climber ascends to an altitude where total pressure is 380 mmHg. Assuming the same fraction of O₂ in air (~21%), what is the approximate alveolar PO₂, and what physiological consequences would you predict?
A classmate says: 'When you mix nitrogen and oxygen in a tank, the gases interact slightly, so the total pressure is a bit less than the sum of their individual pressures.' What's wrong with this statement, and what assumption does Dalton's law make that your classmate is violating?
A patient is breathing supplemental O₂ through a mask, raising the fraction of inspired O₂ from 21% to 50% at sea level. Using Dalton's law, calculate the new PO₂ in inspired air and explain why this intervention helps a hypoxic patient more than it would help a healthy person at extreme altitude.

Related topics

Ideal Gas Law and Real Gas Behavior Density and Specific Gravity Pressure in a Fluid (Pascal's Principle, Hydrostatic) Buoyancy and Archimedes' Principle Continuity Equation (Conservation of Flow)

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