Pressure is force per unit area applied perpendicular to a surface; in AP Chem, gas pressure comes from particle collisions with container walls and connects the ideal gas law (PV = nRT), Dalton's partial pressures, and Le Châtelier's principle for gas-phase equilibria.
Pressure is force per unit area, applied perpendicular to a surface. For gases, that force comes from billions of particles slamming into the container walls every second. More collisions, or harder collisions, means more pressure. That's why squeezing a gas into a smaller volume, adding more moles, or raising the temperature all increase pressure.
In AP Chem, pressure shows up as a measurable macroscopic property tied to particle-level behavior. The ideal gas law (PV = nRT) links it to volume, moles, and temperature (EK 3.4.A.1). In a gas mixture, each gas exerts its own partial pressure independent of the others, and the total pressure is just the sum of the parts (EK 3.4.A.2). Then in Unit 7, changing the pressure of a gas-phase system becomes a 'stress' that knocks an equilibrium out of balance and forces it to shift (EK 7.9.A.1).
Pressure is one of the few concepts that gets tested in at least three different units. In Unit 3, LO 3.4.A asks you to relate pressure to volume, moles, and temperature through PV = nRT and to calculate partial pressures with mole fractions. In Unit 7, LOs 7.9.A and 7.10.A ask you to predict how an equilibrium responds when pressure changes, and to back up that prediction with Q versus K reasoning. Pressure also lurks in Unit 6 (external pressure determines when a liquid boils during phase changes, Topic 6.5) and Unit 5 (compressing a gas raises concentration, which raises reaction rate, EK 5.1.A.3). If you only memorize 'pressure shifts equilibrium toward fewer moles of gas' without understanding the particle-level cause, the trickier exam questions will catch you.
Keep studying AP Chemistry Unit 5
Ideal Gas Law (Unit 3)
PV = nRT is the home base for every pressure calculation. Decreasing volume at constant temperature raises pressure because the same number of particles hit the walls more often. This is Boyle's Law hiding inside the bigger equation.
Dalton's Law of Partial Pressure (Unit 3)
Each gas in a mixture exerts pressure as if it were alone, so P_A = P_total × X_A and the total is the sum of the parts. This idea is the bridge to Unit 7, because Kp and Q for gas reactions are written in partial pressures, not total pressure.
Le Châtelier's Principle (Unit 7)
Shrinking the volume of a gas-phase equilibrium raises every partial pressure, makes Q differ from K, and the system shifts toward the side with fewer moles of gas to bring Q back to K. The Q vs. K explanation (Topic 7.10) earns points; 'the system relieves stress' alone often doesn't.
Energy of Phase Changes (Unit 6)
Boiling happens when a liquid's vapor pressure equals the external pressure pushing down on it. That's why water boils below 100°C at high altitude, where atmospheric pressure is lower.
Pressure questions usually make you do one of three things. First, calculate it. Released FRQs like 2017 Long Q1 (CS₂ + Cl₂) and 2018 Long Q2 (the NO/NO₂ mixture) have asked for partial pressures and mole fractions in closed containers at constant temperature, which is straight PV = nRT and Dalton's Law work. Second, predict equilibrium shifts. MCQs love combination stems like 'which set of stresses most increases PCl₃ in PCl₅ ⇌ PCl₃ + Cl₂?' where you have to weigh a volume change against a temperature change using mole counts and the sign of ΔH°. Third, spot the classic trap. Adding an inert gas like argon at constant volume changes total pressure but not any partial pressure, so Q still equals K and nothing shifts. That exact scenario shows up in multiple-choice questions, and answering it correctly requires the partial-pressure reasoning, not the memorized shortcut.
Total pressure is what a gauge on the container reads; partial pressure is the share contributed by one specific gas, proportional to its mole fraction. Equilibrium (Kp and Q) only cares about partial pressures. That's why adding argon at constant volume raises total pressure but shifts nothing, because no reactant or product's partial pressure changed.
Pressure is force per unit area, and for gases it comes from particle collisions with container walls, so more collisions or more energetic collisions means higher pressure.
The ideal gas law PV = nRT connects pressure to volume, moles, and temperature, and it's your tool for nearly every gas calculation on the exam.
In a mixture, each gas's partial pressure equals its mole fraction times the total pressure, and the partial pressures add up to the total (Dalton's Law).
Decreasing the volume of a gas-phase equilibrium shifts the reaction toward the side with fewer moles of gas, because the shift brings Q back to K.
Adding an inert gas at constant volume does NOT shift equilibrium, since the partial pressures of reactants and products don't change.
A liquid boils when its vapor pressure equals the external pressure, which is why boiling point drops at high altitude.
Pressure is force per unit area applied perpendicular to a surface. For gases it results from particle collisions with container walls, and it links to volume, moles, and temperature through PV = nRT (Topic 3.4).
No, not at constant volume. Argon raises the total pressure but doesn't change the partial pressure of any reactant or product, so Q still equals K and the equilibrium doesn't move. This is one of the most common trap questions in Unit 7.
Pressure (total) is the combined push of all gases in a container; partial pressure is one gas's individual contribution, equal to its mole fraction times the total (P_A = P_total × X_A). Kp and the reaction quotient Q are built from partial pressures, never total pressure.
If pressure increases because volume decreased, the equilibrium shifts toward the side with fewer moles of gas. For 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), compressing the container shifts it right, from 3 moles of gas to 2.
No. Pressure and concentration changes only change Q, and the system shifts to bring Q back to K. Only a temperature change actually changes the value of K (EK 7.10.A.2).