Total pressure is the sum of the partial pressures of every gas in a mixture (P_total = P_A + P_B + P_C + ...), because each ideal gas exerts pressure independently of the others. On the AP Chem exam, it links mole fractions to partial pressures and shows up in both gas law and Kp equilibrium problems.
Total pressure is the combined pressure that all the gases in a container exert together. The big idea behind it (Essential Knowledge 3.4.A.2) is that ideal gases ignore each other. Each gas pushes on the walls as if it were alone in the container, and those individual pushes are called partial pressures. Add them all up and you get the total pressure. That's Dalton's Law of Partial Pressures in one line: P_total = P_A + P_B + P_C + ...
The useful flip side is that each gas's share of the total pressure matches its share of the moles. That's the mole fraction relationship, P_A = P_total × X_A, where X_A = moles of A ÷ total moles. So if N₂ makes up 30% of the moles in a flask, it contributes 30% of the total pressure. Think of total pressure like a group bill at a restaurant. Each gas pays its own share (partial pressure), nobody's order affects anyone else's, and the bill is just the sum.
Total pressure lives in Topic 3.4 (Ideal Gas Law) under learning objective 3.4.A, which asks you to explain the macroscopic properties of a gas mixture using PV = nRT. You can apply the ideal gas law to the whole mixture using total moles, or to one component using its moles alone, and total pressure is what ties those two views together. Then it comes back in Unit 7. Under learning objective 7.1.A, equilibrium for gas-phase reactions is defined by partial pressures becoming constant (EK 7.1.A.2), and equilibrium constants for gases (Kp) are written in terms of partial pressures. A ton of equilibrium problems hand you a total pressure and make you split it into partial pressures before you can do anything else. If you can't move between total pressure, partial pressure, and mole fraction quickly, both Unit 3 and Unit 7 get harder than they need to be.
Partial Pressure (Unit 3)
Total pressure is literally built from partial pressures. Each gas's partial pressure is independent of the others, so the total is just the sum. Most exam problems make you go one direction or the other, splitting a total into parts or adding parts into a total.
Dalton's Law of Partial Pressure (Unit 3)
Dalton's Law is the name for the rule that total pressure equals the sum of partial pressures. It works because ideal gas particles don't attract or repel each other, so each gas behaves as if it has the container to itself.
Ideal Gas Law (Unit 3)
PV = nRT works for the whole mixture if n is the total moles, and the P you get is the total pressure. That's why adding 0.20 mol of He to a rigid container of N₂ raises the total pressure in a way you can calculate directly from n, R, T, and V.
Chemical Equilibrium (Unit 7)
At equilibrium in a gas-phase reaction, the partial pressures of all species stop changing (EK 7.1.A.2), which means total pressure stops changing too. Equilibrium problems often give you a measured total pressure and ask you to back out the partial pressures that go into Kp.
In multiple choice, total pressure questions are usually quick math. You'll add a gas to a rigid container and find the new total pressure, convert between total pressure and mole fraction (like finding the mole fraction of N₂ in a 1.50 atm mixture of H₂, N₂, and CO₂), or remove one gas and predict the new total (remove the CO₂ that contributed half the moles, and the total drops by half). The trap answers usually come from forgetting that each gas's pressure is independent of the others. On free response, total pressure shows up inside gas-phase equilibrium and stoichiometry problems. The 2022 FRQ on methanol decomposition (CH₃OH → CO + 2H₂) is the classic setup, where the mole count changes as the reaction proceeds, so the total pressure in a rigid container changes too, and you reason from a measured total pressure back to partial pressures of individual species. Be ready to write P_total = P_A + P_B + ... and P_A = P_total × X_A from memory and justify when each applies.
Partial pressure is the pressure ONE gas in a mixture exerts on its own; total pressure is the sum of all of them. A pressure gauge on the container reads total pressure, never an individual partial pressure. If a problem says 'the pressure in the flask is 2.5 atm,' that's total pressure, and you need mole fractions (or stoichiometry) to split it into the partial pressures that go into Kp.
Total pressure equals the sum of all partial pressures in a gas mixture: P_total = P_A + P_B + P_C + ... (Dalton's Law).
Each gas's partial pressure is independent of the other gases, so its share of the total pressure equals its mole fraction: P_A = P_total × X_A.
Adding gas to a rigid container at constant temperature raises the total pressure by exactly the new gas's partial pressure; removing a gas drops the total by that gas's contribution.
A pressure gauge measures total pressure, so equilibrium problems often require you to convert a measured total pressure into partial pressures before plugging into Kp.
At equilibrium, the partial pressures of all gaseous species are constant (EK 7.1.A.2), so a constant total pressure in a rigid container is one observable sign that a gas-phase system has reached equilibrium.
Total pressure is the sum of the partial pressures of every gas in a mixture, P_total = P_A + P_B + P_C + ... It appears in Topic 3.4 (Ideal Gas Law) under EK 3.4.A.2 and again in Unit 7 equilibrium, where Kp is built from partial pressures.
No. Partial pressure is what one gas alone contributes, while total pressure is what all the gases contribute combined. A gauge on the container always reads total pressure, and you use mole fractions to split it into partial pressures.
Yes, in a rigid container it does. Adding 0.20 mol of He to a sealed flask adds He's partial pressure (from PV = nRT) on top of whatever was already there, since ideal gases exert pressure independently. The original gas's partial pressure stays the same.
Multiply the total pressure by the gas's mole fraction: P_A = P_total × X_A, where X_A is moles of A divided by total moles. For example, 0.20 mol N₂ out of 0.70 total mol in a 1.50 atm mixture gives X = 0.29 and P_N₂ ≈ 0.43 atm.
Gas-phase equilibrium constants (Kp) are written in terms of partial pressures, but experiments usually measure total pressure. FRQs like the 2022 methanol decomposition question expect you to use stoichiometry and mole fractions to turn a measured total pressure into the partial pressures Kp needs.
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