Adding more reactant usually shifts a chemical equilibrium toward the products because the disturbance makes the mixture inconsistent with its equilibrium ratio. The system then undergoes a net forward reaction, consuming some of the added reactant and producing more product until equilibrium is restored.
This is the concentration effect predicted by Le Chatelier's principle. For a deeper explanation, however, IB Chemistry students should connect the verbal rule to two ideas: the temporary inequality between the forward and reverse rates, and, particularly at HL, the relationship between the reaction quotient, Q, and the equilibrium constant, K.
Le Chatelier's principle explained for concentration changes
Le Chatelier's principle states that when a system at equilibrium is disturbed, the position of equilibrium changes in a direction that partially counteracts the disturbance. In the current IB Chemistry course, this appears within Reactivity 2.3: How far? The extent of chemical change, with the qualitative application of Le Chatelier's principle identified as R2.3.4.
For concentration changes, the basic predictions are:
Change imposedDirection of net reactionResultAdd a reactantToward productsSome added reactant is consumedRemove a reactantToward reactantsMore reactant is producedAdd a productToward reactantsSome added product is consumedRemove a productToward productsMore product is produced
The word partially matters. If a reactant concentration is increased, the system normally consumes only part of the added amount. Its final concentration can therefore remain higher than it was at the original equilibrium.
Students who need broader topic coverage should first review Le Chatelier's principle in IB Chemistry explained. The present article focuses specifically on why adding reactant produces a shift rather than repeating the entire equilibrium topic.
What an equilibrium shift actually means
Consider the general reversible reaction:
aA + bB ⇌ cC + dD
At dynamic equilibrium, the forward and reverse reactions continue, but their rates are equal. As a result, the concentrations of A, B, C, and D remain constant over time, although they are not necessarily equal to one another.
A shift to the right means that there is a temporary net forward reaction:
-
Reactants are consumed overall.
-
Products are formed overall.
-
The forward rate is temporarily greater than the reverse rate.
-
The process continues until the two rates become equal again.
Strictly speaking, an equilibrium itself does not move while remaining at equilibrium. Immediately after reactant is added, the mixture is no longer at equilibrium. Saying that “the equilibrium shifts right” is convenient shorthand for saying that the system undergoes a net forward reaction and establishes a new equilibrium composition containing relatively more products.
The RevisionDojo explanation of dynamic equilibrium develops the distinction between equal rates and constant concentrations.
Why adding reactant causes a forward reaction
There are three useful levels of explanation. Le Chatelier's principle provides the prediction, rate changes describe the particle-level response, and the reaction quotient gives the most precise mathematical justification.
The Le Chatelier explanation
Adding a reactant increases its concentration. The system responds in the direction that consumes that substance, so the forward reaction is favoured and more product forms.
For example:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Adding nitrogen or hydrogen causes a net reaction to the right. Nitrogen and hydrogen are consumed as additional ammonia forms.
This is usually sufficient for a short qualitative question, provided the answer identifies the disturbance, direction, and consequence. Avoid writing only “the system opposes the change,” because that does not explicitly state what changes chemically.
The rate explanation
Before the addition, the forward and reverse rates are equal. Adding a reactant suddenly increases the number of its particles per unit volume, so encounters involving that reactant become more frequent and the forward reaction rate generally increases immediately.
The product concentrations have not changed at that instant, so the reverse rate does not receive the same immediate increase. The forward rate therefore becomes greater than the reverse rate, creating a net conversion of reactants into products.
As reactants are consumed and products accumulate, the forward rate decreases relative to its value just after the addition, while the reverse rate increases. Eventually, the rates become equal again at a new dynamic equilibrium.
This particle explanation is intuitive, but the reaction mechanism can make actual rate laws more complicated than the balanced equation suggests. For equilibrium-direction questions, comparing Q and K is the more general and reliable method.
The reaction quotient explanation
For the general reaction
aA + bB ⇌ cC + dD
the concentration reaction quotient is written as:
Qc = ([C]^c[D]^d) / ([A]^a[B]^b)
It has the same mathematical form as the equilibrium expression. The difference is that Q can be calculated at any moment, whereas K describes the equilibrium value at a specified temperature.
At equilibrium:
Q = K
If more A is added, [A] increases. Because [A]^a is in the denominator, Q immediately decreases, giving:
Q < K
The reaction must proceed to the right to increase Q. It does this by decreasing reactant concentrations and increasing product concentrations until Q = K again. This relationship is the rigorous reason why equilibrium shifts when more reactant is added.
ComparisonMeaningDirection of net reactionQ < KToo much reactant relative to the equilibrium ratioRight, toward productsQ = KEquilibrium compositionNo net reactionQ > KToo much product relative to the equilibrium ratioLeft, toward reactants
This treatment is especially important for IB Chemistry HL, where the current syllabus includes the reaction quotient under R2.3.5. The RevisionDojo R2.3 equilibrium lessons separate the qualitative Le Chatelier treatment from the quantitative use of Q.
Worked example using Q and K
Consider:
H₂(g) + I₂(g) ⇌ 2HI(g)
The reaction quotient is:
Qc = [HI]² / ([H₂][I₂])
Suppose the mixture is initially at equilibrium, so Qc = Kc. If hydrogen is suddenly added at constant volume, [H₂] increases while [I₂] and [HI] are unchanged at that instant.
The denominator becomes larger, so Qc falls below Kc. The forward reaction then occurs overall:
-
[H₂]decreases from its value immediately after addition. -
[I₂]decreases. -
[HI]increases. -
Qcrises until it equalsKcagain.
Notice that the system does not necessarily return [H₂] to its original equilibrium concentration. Some added hydrogen is consumed, but the new equilibrium hydrogen concentration will commonly remain above the original value.
The stoichiometric changes also follow the balanced equation. If the forward shift consumes x mol dm⁻³ of H₂, it consumes x mol dm⁻³ of I₂ and forms 2x mol dm⁻³ of HI.
Why K does not change when reactant is added
A concentration change alters Q, not K, as long as the temperature remains constant. The equilibrium constant depends on temperature, so simply adding reactant does not create a new value of Kc.
This distinction is central to IB chemistry equilibrium questions:
-
Q changes immediately when a relevant concentration changes.
-
The system undergoes a net reaction because
Q ≠ K. -
Concentrations change as the system re-equilibrates.
-
Q returns to K at the new equilibrium.
-
K remains unchanged if temperature has not changed.
Do not write that adding reactant “increases K” or “makes K favour products.” Adding reactant can increase the equilibrium yield of product, but it does not alter the thermodynamic equilibrium constant at that temperature.
Cases where adding a reactant does not give the simple result
The rule “add reactant, shift right” assumes that adding the substance increases an effective concentration or partial pressure appearing in the reaction quotient. Several important qualifications prevent careless application.
Adding a pure solid
Pure solids are omitted from heterogeneous equilibrium expressions because their activity is treated as constant. Adding more of a pure solid therefore does not shift equilibrium, provided some of that solid was already present and the temperature and other conditions remain unchanged.
For example:
CaCO₃(s) ⇌ CaO(s) + CO₂(g)
The equilibrium expression depends on CO₂ but not on the amounts of the two pure solids. Adding more CaCO₃ does not change Q and therefore does not cause a shift. It may allow more material to react if another condition changes, but it does not alter the equilibrium gas pressure by itself.
Adding reactant together with extra solvent
If a reactant solution is added, both the amount of reactant and the total volume may change. The concentration is n/V, so adding more moles does not guarantee that the reactant's concentration increases by the same proportion.
In a careful calculation, determine the new concentrations after mixing and then calculate Q. This is safer than applying a memorized rule based only on the number of moles added.
Adding a gas while changing pressure or volume
Adding a gaseous reactant at constant volume increases its partial pressure and normally gives the expected shift toward products. If the container volume also changes, however, all gaseous partial pressures may be affected.
Calculate Q from the new partial pressures or concentrations if the outcome is not obvious. This prevents concentration and pressure effects from being confused.
The reaction cannot establish equilibrium
Dynamic equilibrium requires conditions that allow both forward and reverse reactions to occur. If products escape from an open system, or if one phase is physically removed, the original closed equilibrium model may no longer apply in the simple way assumed by an exam question.
How to answer this in an IB Chemistry exam
For a standard explanation question, use a short cause-and-effect chain:
-
State the disturbance: the concentration of a reactant increases.
-
State the direction: the system undergoes a net forward reaction, so the equilibrium position shifts toward products.
-
Explain the response: some of the added reactant is consumed and more product forms.
-
Mention K when relevant: the equilibrium constant remains unchanged because temperature is constant.
A strong response could be:
Increasing the concentration of a reactant causes the system to shift toward the products, partially consuming the added reactant. Immediately after the addition, Q is less than K, so a net forward reaction occurs until Q equals K again. K remains constant because the temperature has not changed.
At SL, a clear Le Chatelier explanation may be enough for a qualitative prompt. At HL, or whenever Q and K are supplied, compare them explicitly rather than relying only on the verbal rule.
For focused practice, use the IB Chemistry equilibrium Questionbank and check whether each response identifies the change, direction, and chemical consequence. The Le Chatelier flashcards are useful for recalling the basic patterns, while Jojo AI can help identify missing reasoning in a written explanation.
Common mistakes to avoid
MistakeWhy it is incorrectBetter statement“The reaction shifts right permanently”The net reaction lasts only until equilibrium is restoredA net forward reaction occurs until a new equilibrium is established“All the added reactant is used up”The disturbance is only partially counteractedSome of the added reactant is consumed“K increases”Concentration changes affect Q, not KK is unchanged at constant temperature“Equilibrium means equal concentrations”Equilibrium requires equal forward and reverse ratesConcentrations are constant, but not necessarily equal“Adding any reactant shifts equilibrium”Pure solids and pure liquids do not normally affect QCheck whether the species appears in the equilibrium expression“A catalyst shifts equilibrium right”A catalyst accelerates both directions and does not change KA catalyst helps equilibrium be reached faster
The RevisionDojo guide to common equilibrium mistakes provides further exam-focused checks covering equilibrium expressions, pressure, temperature, and ICE tables.
Conclusion
Equilibrium shifts when more reactant is added because the new composition no longer satisfies the equilibrium relationship. The added reactant lowers Q relative to K, creating a net forward reaction that consumes some reactant and forms additional product until Q = K again.
Le Chatelier's principle predicts the rightward shift, while rate changes and the Q-versus-K comparison explain why it happens. Remember that the response is partial, K remains constant at constant temperature, and adding a pure solid does not normally change the equilibrium position. RevisionDojo's equilibrium notes, flashcards, Questionbank, and Jojo AI feedback can help you practise expressing this reasoning in the concise form expected in IB examinations.
Sources and referenced URLs
-
International Baccalaureate: Chemistry in the Diploma Programme
-
International Baccalaureate DP Chemistry subject brief, first assessment 2025
-
OpenStax Chemistry 2e: Shifting equilibria and Le Chatelier's principle
-
RevisionDojo: Le Chatelier's principle in IB Chemistry explained
-
RevisionDojo: Dynamic equilibrium explained for IB Chemistry
