Equilibrium thus far has considered only reactions going straight to equilibrium and then staying there. However, it doesn't have to be that way. We can actually control what happens in a chemical system...👀
What if we wanted to say, shift an equilibrium towards the products or towards the reactants? How does equilibrium shift based on certain stressors that are added to a system? This section will address that very question by introducing a new rule: Le Châtelier’s Principle.
What Is Le Châtelier’s Principle?
Le Châtelier’s Principle states one key rule about how a system in equilibrium will react to an external pressure: “if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change to reestablish an equilibrium” (LibreTexts). Let’s break down each part of this.
Dynamic Equilibrium and Stressors
We begin by establishing that we are in a dynamic equilibrium. This just means that we have some reaction that is at equilibrium. We call it dynamic equilibrium because equilibrium does not mean the reaction has stopped but rather that the concentrations of reactants and products have not changed. You may be wondering what other type of equilibrium there is. There is also something called static equilibrium and it's basically a state of balance in which the system is not moving. No need to worry about that, try to focus on dynamic equilibrium!
Next, our dynamic equilibrium is disturbed by “changing the conditions”. We’ll cover some specific disturbances in the next section, but generally, this phrase refers to any external change that comes to the system from outside. This could mean adding more of something, increasing temperature, changing pressures, etc. Remember when we discussed how external factors can change the kinetics of a reaction? Well...they can affect where a reaction stands with regard to equilibrium.
Finally, and this is the part where Le Châtelier’s Principle answers our question, the equilibrium will shift to counteract the change and re-establish equilibrium. Think of Le Châtelier’s principle as a seesaw that always wants to remain in balance. If you add a 50-pound weight to one side, Le Châtelier’s principle will respond by counteracting that change and shifting some of that 50lbs to the other side.
Factors That Influence Equilibrium
There are several factors that can influence equilibrium in a chemical or physical system and you should understand how Le Châtelier’s applies.
Concentration
Changes in concentration are the clearest way in which equilibrium can shift. The concentrations of the reactants and products are constant at equilibrium, so what happens if external stress adds or removes substances? Note that the K value of the reaction is not going to change, just the concentrations at equilibrium. This is referred to as inducing new stress to the system, and based on Le Châtelier’s Principle, we can figure out which direction the reaction will shift to re-establish equilibrium.
If we add more reactants to the system, the system will respond by counteracting that change and creating more products to return to equilibrium. Similarly, if we add products to a system the system will respond and create more reactants. Changes in concentration as stressors are pretty easy to compensate, as you can consume the substances there are more of and create the substances there are less of.
Let's consider the following chemical equation: Fe³⁺ + SCN⁻ ⇌ FeSCN²⁺. Imagine the system is at equilibrium, and we add a compound containing the SCN⁻ ion to the system, increasing its concentration. That excess ion is going to be consumed to create more products and re-establish the equilibrium. We can see this in the graph below.
Notice that there is a sharp spike in the concentration of SCN⁻, indicating that its addition is the stress the system responds to. Once the [SCN⁻] is added, we see a decrease in [Fe³⁺] and [SCN⁻] as responses to the stress: the system responds by creating more [FeSCN]²⁺.
Another important part of how concentration impacts equilibrium is adding a compound that will react with a reactant or product to take it out of the system. For example, adding a compound containing hydroxide will cause the hydroxide to react to form Fe(OH)₃, taking the Fe³⁺ ion out of the system. Based on Le Châtelier’s principle, the system will respond by increasing the amount of Fe³⁺ and SCN⁻ present in the system and decreasing the concentration of [FeSCN]²⁺.
Temperature
Temperature changes, like changes in concentration, have an impact on equilibrium. In this case, the change in equilibrium depends on whether or not our reaction is exothermic (heat releasing) or endothermic (heat absorbing). Remember from unit six that we can identify whether a reaction is exothermic or endothermic by observing its change in standard enthalpy, or ΔH°.
If ΔH° < 0, or negative, our reaction is exothermic, and if ΔH° > 0, or positive, our reaction is endothermic. When referring to Le Châtelier’s Principle, we can think of heat as either a reactant or product and then refer to concentration for similar conclusions for temperature.
For example, consider the reaction N₂ + 3H₂ ⇌ 2NH₃. This reaction has a ΔH° of -92 kJ/mol. Since ΔH° is negative, our reaction is exothermic and heat is released into the surroundings, implying that heat can be thought of as a product. This means that if the temperature increases, our equilibrium will shift to produce more reactants. We're basically fueling the reaction that takes in heat, which in this case, is the reverse reaction.
The opposite would occur if the reaction was endothermic and the temperature was a reactant. With temperature, try to think about which direction of the reaction you are "fueling;" that is the direction the reaction will shift.
Pressure
When dealing with a system involving gases, changes in pressure (and therefore the volume, remember Boyle’s Law!) will change concentrations of species in the system.
In general, if the pressure on a system increases or the volume decreases, the equilibrium will shift towards the side with fewer moles of gas and vice versa. If there is no difference in moles of gas, there will be no change either way.
When we talk about moles of gas, we’re talking about the coefficients in the reaction.
For example, consider the reaction A + B ⇌ 2C + 3D where A B C and D are gasses. On the left, there are 2 total moles of gas because 1 + 1 = 2. On the right, there are 5 total moles because 2 + 3 = 5. Therefore, if the pressure were to increase on the system, the equilibrium would shift left and products would be converted into more reactants.
However, there is one caveat to this rule that has to do with inert gases. An inert gas is a gas that does not react when added to a mixture of chemicals. If an inert gas is added or removed from a vessel, there is no impact on equilibrium. For example, if Helium gas was pumped in, there would be no impact as far as equilibrium is concerned.
Summary of Le Châtelier’s Principle
We just went over so many rules, so here is a nice table that can break them down a bit:
| Stress | Shift | Explanation |
|---|---|---|
| Increase the concentration of a substance | Away from the substance | Extra concentration needs to be used up |
| Decrease the concentration of a substance | Towards the substance | Need to produce more of the substance that was removed to make up for the loss |
| Increase the pressure of the system | Toward fewer moles of gas | Boyle's law: pressure increase = volume decrease |
| Decrease the pressure of the system | Toward more moles of gas | Boyle's law: pressure decrease = volume increase |
| Increase the temperature of the system | Away from the heat, favoring the endothermic reaction | Extra heat has to be used up to fuel the reaction |
| Decrease the temperature of the system | Toward the heat, favoring the exothermic reaction | More heat needs to be produced to make up for the loss |
| Adding a catalyst | --no shift-- | The rates of both the forward and reverse reactions are increased by the same amount. Adding a catalyst only affects kinetics, not equilibrium. |
Vocabulary
The following words are mentioned explicitly in the College Board Course and Exam Description for this topic.
| Term | Definition |
|---|---|
| chemical species | A distinct chemical entity such as an atom, molecule, or ion that participates in a chemical reaction. |
| dilution | The process of decreasing the concentration of a solute in a solution by adding solvent, which can shift equilibrium position. |
| equilibrium | The state in which the forward and reverse reaction rates are equal, resulting in constant concentrations or partial pressures of reactants and products. |
| external stress | A change applied to a system at equilibrium, such as addition or removal of a chemical species, temperature change, pressure change, or dilution. |
| Le Châtelier's principle | A principle stating that when a system at equilibrium is disturbed, the system shifts to counteract the disturbance and re-establish equilibrium. |
| pH | A logarithmic scale used to express the concentration of hydronium ions in a solution, calculated as −log[H3O+]. |
| pressure | The force exerted by gas molecules; changes in pressure of a gas-phase system can shift the equilibrium position. |
| temperature | A factor that influences reaction rate by affecting the kinetic energy and collision frequency of reactant molecules. |
| volume | The space occupied by a system; changes in volume of a gas-phase system can shift equilibrium position. |
Frequently Asked Questions
What is Le Châtelier's principle and how does it work?
Le Châtelier’s principle says that if an equilibrium system is stressed (change in concentration, pressure/volume, temperature, or dilution), the system shifts to partially counteract that stress and restore equilibrium. Practically: add a reactant and the reaction shifts toward products; remove a product and it shifts toward products; increase pressure (by decreasing volume) for gas reactions and the system shifts to the side with fewer moles of gas; change temperature and treat heat as a reactant (for endothermic reactions, adding heat shifts right; for exothermic, adding heat shifts left). Use Q versus K to predict direction quantitatively: if Q < K, reaction goes forward; if Q > K, it goes reverse. Le Châtelier also predicts measurable changes (pH, color, temperature)—e.g., common-ion or dilution effects change [H+] and pH. This is exactly what AP learning objective 7.9.A tests—practice applying these stresses (see the Topic 7.9 study guide for examples) (https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN). For more unit review and lots of practice problems, check the Unit 7 page (https://library.fiveable.me/ap-chemistry/unit-7) and the practice bank (https://library.fiveable.me/practice/ap-chemistry).
Why does adding more reactants make the equilibrium shift to the right?
When you add more reactant, you’ve disturbed the equilibrium by increasing the concentrations on the reactant side. Le Châtelier’s principle says the system will respond to reduce that stress—here by consuming some of the extra reactant. Practically that means the forward reaction speeds up relative to the reverse until a new equilibrium is reached. You can also see it with Q and K: adding reactant lowers Q (makes Q < K), so the reaction must proceed to the right (produce products) until Q = K again. This is exactly the kind of concentration stress Le Châtelier’s principle and Learning Objective 7.9.A predict (see the Topic 7.9 study guide for a quick review: https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN). For extra practice on Q vs K problems, try Fiveable’s AP Chem practice set (https://library.fiveable.me/practice/ap-chemistry).
I'm confused about Le Châtelier's principle - can someone explain it in simple terms?
Think of a chemical equilibrium like a party where people can move between two rooms (reactants ⇌ products). Le Châtelier’s principle says: if you stress the party, the system shifts to reduce that stress. Simple examples you should memorize for AP (LO 7.9.A): - Add/remove a reactant or product: equilibrium shifts away from what you add and toward what you remove (concentration stress). - Change temperature: treat heat as a reactant (endothermic) or product (exothermic). Only temperature changes K. - Change pressure/volume for gases: increasing pressure (decreasing volume) shifts toward the side with fewer moles of gas. - Dilution/common-ion: adding solvent or a common ion shifts accordingly (affects concentrations, pH, solubility). Quick tip: compare Q (reaction quotient) to K. If Q < K, reaction moves forward (toward products); if Q > K, it goes backward (toward reactants). Le Châtelier helps predict observable changes (pH, color, temperature). For practice and AP-style questions, use the Topic 7.9 study guide (https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN), the Unit 7 overview (https://library.fiveable.me/ap-chemistry/unit-7), and lots of practice problems (https://library.fiveable.me/practice/ap-chemistry).
What happens to equilibrium when you change the temperature of a reaction?
Temperature changes are treated like adding or removing heat. Use Le Châtelier: if a reaction is exothermic (heat is a product, ΔH < 0), adding heat (raising T) shifts equilibrium toward reactants; removing heat shifts it toward products. If the reaction is endothermic (heat is a reactant, ΔH > 0), adding heat shifts equilibrium toward products; cooling shifts it toward reactants. Unlike concentration or pressure changes, a temperature change also changes the equilibrium constant K. For an exothermic reaction, increasing T decreases K; for an endothermic reaction, increasing T increases K. So both the direction of the shift and the numeric value of K depend on ΔH. On the AP exam you should state the shift using Le Châtelier’s principle and note whether K increases or decreases (CED 7.9.A.1). Want a quick review or practice problems? See the Topic 7.9 study guide (https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN), the Unit 7 overview (https://library.fiveable.me/ap-chemistry/unit-7) and extra practice (https://library.fiveable.me/practice/ap-chemistry).
How do I predict which way the equilibrium will shift when pressure changes?
If pressure changes, look only at stresses that change partial pressures of gases. For a volume decrease (total pressure up) the equilibrium shifts toward the side with fewer moles of gas; for a volume increase (pressure down) it shifts toward the side with more gas molecules. That prediction comes from Le Châtelier’s principle and the pressure–volume relationship (PV work) in gaseous equilibria. Remember K (at constant T) doesn’t change—only Q vs. K determines shift. Adding an inert gas at constant volume has no effect on the equilibrium (partial pressures of reactants/products unchanged); adding an inert gas at constant pressure changes volume and can shift the equilibrium. Temperature changes do change K, so treat those separately. This is an AP-style skill (CED 7.9.A.1); practice predicting shifts with gas equations and counting gas moles. For a quick refresher, see the Topic 7.9 study guide (https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN) and try practice questions (https://library.fiveable.me/practice/ap-chemistry).
What's the difference between adding a reactant and removing a product in terms of equilibrium?
They’re effectively the same kind of stress: both make Q differ from K so the system shifts to re-establish equilibrium. - Adding a reactant increases its concentration → Q (reaction quotient) becomes smaller than K for a reaction written as products/reactants, so the equilibrium shifts to the right (makes more products) to use up the added reactant. - Removing a product decreases its concentration → Q becomes smaller than K too, so the equilibrium also shifts to the right to replace the removed product. Key idea: what matters is how Q compares to K after the stress. If Q < K, the reaction shifts toward products; if Q > K, it shifts toward reactants. Use Le Châtelier’s principle and the Q vs K check to predict direction (CED 7.9.A.1). For more examples and practice, see the Topic 7.9 study guide (https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN) and AP practice problems (https://library.fiveable.me/practice/ap-chemistry).
Why does diluting a solution affect the equilibrium position?
Diluting a solution lowers the concentrations of all dissolved species. Because the reaction quotient Q uses those concentrations, dilution usually changes Q but not the equilibrium constant K (unless temperature changes). Le Châtelier’s principle says the system will shift to oppose that stress—it shifts in the direction that moves Q back toward K. Example: A + B ⇌ C with K = 1. If [A] = [B] = 1 M and [C] = 0.5 M, Q = (0.5)/(1·1) = 0.5 < K, so the equilibrium lies to the right. If you dilute everything by half, Q becomes 0.25, still < K, so the reaction shifts further right (makes more C) to restore K. For equilibria involving H+ (or common ions), dilution changes pH and can shift acid/base equilibria accordingly. This idea is part of AP CED Topic 7.9 (use Q vs K and Le Châtelier to predict shifts). For a quick refresher, see the Topic 7.9 study guide (https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN) and more Unit 7 review (https://library.fiveable.me/ap-chemistry/unit-7). Practice lots of Q vs K problems at (https://library.fiveable.me/practice/ap-chemistry).
How does Le Châtelier's principle help predict changes in pH during a reaction?
Le Châtelier’s principle predicts pH changes by telling you which way an acid–base equilibrium shifts when you stress it. For example: - Add H+ (strong acid): equilibrium of HA ⇌ H+ + A− shifts left, suppressing dissociation (common-ion effect) → [H+] from the added acid dominates, but the weak acid contributes less (pH stays lower than before). - Add OH− (base): OH− removes H+, so equilibrium shifts right to make more H+ (HA dissociates) → pH rises less than if no equilibrium were present. - Dilution: lowers all concentrations, so equilibria shift toward the side with more particles (weak acid dissociation often increases) → pH moves toward that expected from greater dissociation. - Temperature change: if dissociation is endo/exothermic, raising T shifts K (and thus pH); remember K only changes with T. Use Q vs K to predict direction and remember the CED explicitly says you should predict measurable changes like pH (Topic 7.9.A.2). For more examples and practice, see the Topic 7.9 study guide (https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN), the Unit 7 overview (https://library.fiveable.me/ap-chemistry/unit-7), and lots of practice problems (https://library.fiveable.me/practice/ap-chemistry).
What happens to the color of a solution when equilibrium shifts and why?
When an equilibrium shifts, the solution’s color changes if the different species in equilibrium absorb different wavelengths of light. Le Châtelier’s principle predicts which direction the equilibrium moves when you add/remove a reactant, change temperature, dilute, or change pressure; that change alters the concentrations of colored species, so the observed color shifts. Classic AP examples: chromate ⇌ dichromate (yellow ⇌ orange) and the Co2+/CoCl4 2– complex (pink ⇌ blue). If you add something that pushes equilibrium toward the colored product, that color intensifies; if it shifts away, that color fades. Temperature changes matter if the equilibrium is endo- or exothermic (heat acts like a reactant/product). This is exactly the kind of measurable property AP wants you to predict (CED 7.9.A.2). For specific examples and practice, see the Topic 7.9 study guide (https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN), the Unit 7 overview (https://library.fiveable.me/ap-chemistry/unit-7), and more practice problems (https://library.fiveable.me/practice/ap-chemistry).
I don't understand why increasing volume shifts equilibrium toward more gas molecules - can someone help?
Think of pressure (or volume) change as a stress. If you increase the container volume, the total pressure and each gas’s partial pressure drop. Le Châtelier’s principle says the system shifts to oppose that change—it will favor the side with more moles of gas because producing more gas molecules raises the pressure back up. You can also see it with Q and K: increasing V lowers all partial pressures, so the reaction quotient Q (which depends on those partial pressures) shifts relative to K. If the side with more gas moles is favored, the shift continues until Q = K again. Example: N2(g) + 3H2(g) ⇌ 2NH3(g). Increasing V (lowering P) shifts the equilibrium toward the left or right depending on which side has more gas moles (here, left has 4 mol gas vs. 2 mol on right), so it shifts left (toward more gas moles). For the AP exam objective 7.9.A.1 and more practice, see the Topic 7.9 study guide (https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN) and lots of practice questions (https://library.fiveable.me/practice/ap-chemistry).
How do you use Le Châtelier's principle to predict temperature changes in exothermic vs endothermic reactions?
Think of heat as a chemical species. For an exothermic reaction (ΔH < 0), write heat on the product side; for an endothermic reaction (ΔH > 0), put heat on the reactant side. Le Châtelier’s principle says the system shifts to oppose a stress: - Exothermic (heat = product): raising T = add “product” → equilibrium shifts left (toward reactants); lowering T → shifts right (toward products). K decreases with ↑T. - Endothermic (heat = reactant): raising T = add “reactant” → shifts right (toward products); lowering T → shifts left. K increases with ↑T. On the AP, you’ll often be asked to predict the direction of shift or how measurable properties (temperature, color, pH) change (CED 7.9.A.1–7.9.A.2). Remember: Le Châtelier predicts the shift direction, and thermodynamics (ΔG = ΔH − TΔS) explains why K changes with T. For a quick refresher, see the Topic 7.9 study guide (https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN) and more Unit 7 review (https://library.fiveable.me/ap-chemistry/unit-7).
What are all the different types of stresses that can affect chemical equilibrium?
Stresses that can shift an equilibrium (Le Châtelier’s principle)—all AP-relevant types you should know: - Addition or removal of a chemical species (changing concentrations)—including common-ion effects. - Dilution (adding solvent)—lowers concentrations, can shift equilibria. - Change in temperature—note this changes K (heat acts like a reactant for endothermic, product for exothermic). - Change in pressure/volume for gas-phase systems—changing total or partial pressures shifts equilibrium (use partial pressures and Q vs K). - Changes that alter pH (adding acid/base) for equilibria involving H+ or OH− (shifts species and can change color for indicator-type systems). - Changes in particle identity/phase (adding/removing solids usually doesn’t change K but can affect amounts). On the AP exam you should use Q vs K to predict direction and remember only temperature changes K (CED 7.9.A.1–7.9.A.2). For examples and practice, see the Topic 7.9 study guide (https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN) and practice problems (https://library.fiveable.me/practice/ap-chemistry).
Why does removing a product make more product form - isn't that backwards?
Think of equilibrium as a balance: for A + B ⇌ C, at equilibrium Q = K. If you remove product C, you lower the instantaneous Q (Q < K). Le Châtelier’s principle says the system will shift to oppose that stress—here it shifts right (making more C) until Q = K again. It’s not backwards: removing C creates a concentration imbalance that the forward reaction fixes by producing C. Use the reaction-quotient idea on the AP exam (LO 7.9.A): compare Q and K to predict the direction (Q < K → forward; Q > K → reverse). If temperature, pressure, or volume changes, treat those as stresses too (see Topic 7.9 study guide for examples) (https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN). For extra practice with these concepts, try the AP problems at Fiveable (https://library.fiveable.me/practice/ap-chemistry).
How do I know if pressure changes will even affect my equilibrium reaction?
Pressure only matters for equilibria that involve gases. If the reaction has a different number of moles of gas on each side (Δngas ≠ 0), changing the volume/pressure will change partial pressures and shift the equilibrium (Le Châtelier’s principle / CED 7.9.A.1). Example: N2(g)+3H2(g) ⇌ 2NH3(g) has Δn = 2 − 4 = −2; increasing pressure (decreasing volume) favors the side with fewer moles (NH3). If Δngas = 0 (same total gas moles both sides), a pressure/volume change won’t shift the equilibrium (K and Q based on partial pressures scale the same). Adding an inert gas at constant volume doesn’t change partial pressures so it won’t shift equilibrium; adding an inert gas at constant pressure changes volume and can shift it. For quick practice and examples see the Topic 7.9 study guide (https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN) and more problems at the Unit 7 page (https://library.fiveable.me/ap-chemistry/unit-7) or practice bank (https://library.fiveable.me/practice/ap-chemistry).
Can Le Châtelier's principle predict exactly how much the equilibrium will shift or just the direction?
Short answer: Le Châtelier’s principle tells you only the direction the equilibrium will shift when a stress is applied (add/remove a species, change pressure/volume, dilute, or change temperature). It’s a qualitative rule in the AP CED (7.9.A.1–7.9.A.2)—useful for predicting increases/decreases in concentration, pH, color, or temperature, but not the exact new amounts. To find how much the system shifts (the quantitative change), you need the equilibrium constant K and an ICE table (or use Q vs K and solve for equilibrium concentrations/pressures). Note: temperature changes alter K (so you must recalculate K for quantitative predictions), whereas adding/removing species or changing V/pressure does not change K. For practice using both Le Châtelier qualitatively and ICE/K quantitatively, see the Topic 7.9 study guide (https://library.fiveable.me/ap-chemistry/unit-7/intro-le-chateliers-principle/study-guide/ST8UE6kcdhi0hkA5bSkN) and try AP-style problems at (https://library.fiveable.me/practice/ap-chemistry).