Partial pressure is the pressure a single gas in a mixture would exert if it occupied the container alone. On AP Chem, you find it with Dalton's Law (P_A = P_total × X_A) in Unit 3 and plug it into Kp expressions to solve equilibrium problems in Unit 7.
Partial pressure is the pressure that one specific gas in a mixture contributes, as if it were the only gas in the container. Because ideal gases don't interact with each other, each gas behaves independently. The pressure of gas A depends only on how many moles of A are present, not on what else is floating around. That's why partial pressure is proportional to mole fraction: P_A = P_total × X_A, where X_A is moles of A divided by total moles.
Add up every gas's partial pressure and you get the total pressure of the mixture (P_total = P_A + P_B + P_C + ...). That's Dalton's Law of Partial Pressures, straight from the Unit 3 essential knowledge (3.4.A.2). Think of total pressure like a pizza bill split by how many slices each person ate. Each gas's share of the pressure matches its share of the moles. The concept shows up first in Unit 3 with the ideal gas law, then becomes the workhorse of Unit 7, where equilibrium constants for gas-phase reactions (Kp) are written in terms of partial pressures instead of molar concentrations.
Partial pressure lives in two places on the AP Chem exam. In Unit 3, learning objective 3.4.A asks you to relate the macroscopic properties of gas mixtures using the ideal gas law, and the essential knowledge spells out both Dalton's Law equations you need. In Unit 7, partial pressure becomes the currency of equilibrium. LO 7.4.A has you calculate Kp from measured partial pressures at equilibrium, and LO 7.7.A flips it around, asking you to predict equilibrium partial pressures from initial conditions and K. Even the definition of equilibrium itself uses the term. Per 7.1.A.2, a system is at equilibrium when the concentrations or partial pressures of all species stay constant over time. If you can't manipulate partial pressures, roughly two units' worth of quantitative questions are locked.
Keep studying AP Chemistry Unit 3
Dalton's Law of Partial Pressure (Unit 3)
Dalton's Law is the rulebook for partial pressures. Each gas's partial pressure equals total pressure times its mole fraction, and all the partial pressures sum to the total. It works because ideal gases ignore each other, so each one pushes on the walls as if it were alone.
Chemical Equilibrium and Kp (Unit 7)
For gas-phase reactions, you write the equilibrium constant Kp using partial pressures instead of concentrations. Products' partial pressures go on top, reactants' on the bottom, each raised to its coefficient. A Kp problem is just an equilibrium constant expression fed with pressures, so the math feels identical to Kc problems.
Ideal Gas Law (Unit 3)
PV = nRT is the bridge between moles and pressure. Since pressure is proportional to moles at fixed V and T, you can convert between mol fractions and partial pressures, or between concentrations (n/V) and pressures. This is the link that ties Kc and Kp together.
Concentration Changes Over Time (Unit 5)
Kinetics experiments on gases often track partial pressure instead of concentration, because pressure is what a sensor actually measures. Since P is proportional to n/V, a plot of ln(P) vs. time being linear tells you the reaction is first order, just like ln[A] would.
Multiple-choice questions hand you equilibrium partial pressures and ask for Kp. A classic stem gives you 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) with P values like 0.350 atm, 0.175 atm, and 0.700 atm, and you build (P_SO₃)²/[(P_SO₂)²(P_O₂)] and crunch it. Watch the coefficients; squaring is where most points die. Conceptual MCQs test whether you know that constant partial pressures in a closed system signal dynamic equilibrium, not a stopped reaction. On the free-response side, partial pressure is a regular. The 2017 long FRQ (N₂ + O₂ ⇌ NO), the 2022 FRQ on methanol decomposition, and the 2025 and 2026 FRQs on P₄ ⇌ P₂ all involve writing or using Kp expressions with partial pressures. Expect to set up an ICE-style table in atm, compare Q to K to predict which direction the reaction shifts, and solve for an unknown equilibrium pressure.
Total pressure is what the whole mixture exerts on the container; partial pressure is one gas's individual contribution. They're connected by Dalton's Law (total = sum of partials), but only partial pressures go into a Kp expression. A common error is plugging P_total into Kp, or forgetting that adding an inert gas raises total pressure without changing any reacting gas's partial pressure, so the equilibrium doesn't shift.
Partial pressure is the pressure one gas in a mixture would exert if it were alone in the container, and it depends only on that gas's moles, temperature, and volume.
Dalton's Law gives you two equations: P_A = P_total × X_A, and P_total is the sum of all the partial pressures.
Kp expressions use equilibrium partial pressures the same way Kc expressions use concentrations, with each pressure raised to its coefficient from the balanced equation.
A gas-phase system is at equilibrium when the partial pressures of all species stop changing, even though the forward and reverse reactions are still running.
Adding an inert gas at constant volume raises total pressure but leaves every partial pressure unchanged, so it does not shift the equilibrium.
Because pressure is proportional to n/V for an ideal gas, you can treat partial pressure like a stand-in for concentration in both kinetics plots and equilibrium math.
Partial pressure is the pressure a single gas in a mixture exerts on its own, equal to the total pressure times that gas's mole fraction (P_A = P_total × X_A). It comes from Dalton's Law in Topic 3.4 and is the basis of Kp calculations in Unit 7.
No. Total pressure is the sum of every gas's partial pressure in the mixture. If a flask holds 0.350 atm SO₂, 0.175 atm O₂, and 0.700 atm SO₃, the total pressure is 1.225 atm, but only the individual partial pressures go into the Kp expression.
Kp uses equilibrium partial pressures (in atm) and Kc uses molar concentrations (mol/L). The expressions look identical in structure; you just feed them different units. The AP exam tells you which one to calculate, so match your inputs to the constant asked for.
Not at constant volume. The inert gas raises total pressure, but each reacting gas's partial pressure stays the same because its moles and volume haven't changed. Since Q still equals K, the equilibrium does not shift.
No. Constant partial pressures in a closed system mean the system reached dynamic equilibrium (7.1.A.2). The forward and reverse reactions are both still occurring, just at equal rates, so there's no net change to observe.