Course: CHEM 106, Chemistry (NYU) | Source: Lab 9, "Reactants and Products" online simulation Tags: chemical equilibrium, dynamic equilibrium, equilibrium constant, K, Kp, rate constants, Le Chatelier, ICE table, reaction quotient
Difficulty: Introductory to Intermediate
Prerequisites: Basic stoichiometry, mole ratios, concentration vs. pressure expressions. Familiarity with rate laws and the idea that reactions can run in both directions.
Chemical equilibrium is the point where a forward reaction and its reverse run at the same rate, so the concentrations of reactants and products stop changing on a macroscopic level. This lab uses a simulation to let you adjust initial concentrations and rate constants, then watch how the system settles into equilibrium and how the equilibrium constant K relates to the ratio of rate constants. Understanding equilibrium is foundational for acid-base chemistry, solubility, thermodynamics, and virtually every reaction system you will meet from here on.
At equilibrium, reactions have not stopped: both forward and reverse processes continue at equal rates, so net concentrations hold steady. The equilibrium constant K depends on temperature and on the ratio of forward to reverse rate constants, not on starting concentrations. Changing temperature changes the rate constants and therefore changes K.
Chemical equilibrium
The state in a reversible reaction where the forward and reverse reaction rates are equal, so the concentrations of all species remain constant over time. In simple terms, everything is still reacting, but the amounts stop changing because the two directions balance out.
Dynamic equilibrium
Equilibrium in which both the forward and reverse reactions continue to occur. "Dynamic" distinguishes it from a static situation where nothing happens at all. Think of it as two equally matched tug-of-war teams: both are pulling, but the rope does not move.
Equilibrium constant (K)
A numerical value expressing the ratio of product concentrations to reactant concentrations at equilibrium, each raised to the power of its stoichiometric coefficient. A large K means products are favoured; a small K means reactants are favoured. K depends only on temperature for a given reaction.
Kp (equilibrium constant in terms of partial pressures)
The equilibrium constant written using partial pressures instead of molar concentrations. Used for gas-phase reactions. For the reaction H₂(g) + I₂(g) → 2 HI(g), Kp = (P_HI)² / (P_H₂ · P_I₂).
Forward rate constant (k_f)
The rate constant for the reaction proceeding from reactants to products. A larger k_f means the forward reaction is intrinsically faster.
Reverse rate constant (k_r)
The rate constant for the reaction proceeding from products back to reactants.
Relationship K = k_f / k_r
At equilibrium the forward and reverse rates are equal: rate_f = rate_r, so k_f[reactants] = k_r[products]. Rearranging gives K = k_f / k_r. This is the direct link between kinetics and equilibrium.
Rate constant
A proportionality constant in a rate law that depends on temperature (via the Arrhenius equation) but not on concentration. You cannot "set" it in a real lab the way you can in a simulation; you change it by changing temperature.
Le Chatelier's principle
If a system at equilibrium is disturbed (by changing concentration, pressure, or temperature), it shifts to partially counteract that disturbance. Useful for predicting the direction of change, though not explicitly tested in this lab simulation.
Both forward and reverse reactions proceed simultaneously. Macroscopic concentrations are constant, but molecules are still converting back and forth.
"At equilibrium nothing is happening" is false. The correct statement is that the net change in concentration is zero because the two rates match.
For a general reaction aA + bB ⇌ cC + dD, the equilibrium expression is K = [C]^c [D]^d / [A]^a [B]^b.
K is dimensionless in the thermodynamic convention but may carry units in general chemistry courses depending on how it is defined.
K depends only on temperature. Changing initial concentrations does not change K; the system simply adjusts its equilibrium position to satisfy the same K.
At equilibrium: rate_forward = rate_reverse, so k_f [reactants] = k_r [products].
Rearranging: K = k_f / k_r.
In the simulation, when k_f = 1 and k_r = 0, the reaction goes to completion (all reactants become products, K is undefined because the denominator is zero).
When k_f = 0.3 and k_r = 0.1, the ratio k_f/k_r = 3. The experimental K values from different trials (1.36, 2.56, 3.82, 3.33) cluster around 3, with variation from the simulation's stochastic behaviour.
Trial 1 (k_f = 1, k_r = 0): Reaction goes to completion. [R] and [Y] drop to 0, [G] and [B] reach 100. K is undefined (division by zero). This is not a true equilibrium; it is an irreversible reaction.
Trial 2 (k_f = 0.3, k_r = 0.1, starting from reactants only): Equilibrium reached with [R] = [Y] ≈ 46, [G] = [B] ≈ 54. K ≈ 1.36.
Trial 3 (starting from products only, same rate constants): Equilibrium reached with [R] = [Y] ≈ 38, [G] = [B] ≈ 62. K ≈ 2.56.
Trial 4 (starting from equal amounts of all four species): Equilibrium at [R] = [Y] ≈ 34, [G] = [B] ≈ 66. K ≈ 3.82.
Trial 5 (asymmetric starting amounts): Equilibrium at [R] ≈ 10, [Y] ≈ 53, [G] ≈ 27, [B] ≈ 68. K ≈ 3.33.
Key observation: K values vary slightly across trials because the simulation uses a stochastic (random) model, but the ratio k_f/k_r = 3 is the theoretical K. The starting concentrations differ, the equilibrium concentrations differ, but K converges toward the same value.
For H₂(g) + I₂(g) → 2 HI(g), given equilibrium partial pressures P_H₂ = 1.40 × 10⁻² atm, P_I₂ = 3.50 × 10⁻³ atm, P_HI = 2.60 × 10⁻¹ atm:
Kp = (P_HI)² / (P_H₂ × P_I₂)
Kp = (0.260)² / (0.0140 × 0.00350)
Kp = 0.0676 / 4.90 × 10⁻⁵
Kp ≈ 1380
A large Kp tells you products (HI) are heavily favoured at this temperature, which is consistent with the lab's note that there are more products than reactants at room-temperature equilibrium.
Rate constants depend on temperature (Arrhenius equation: k = A e^(−Ea/RT)).
Raising temperature increases both k_f and k_r, but by different amounts depending on the activation energies of the forward and reverse reactions.
Because K = k_f / k_r, changing temperature changes K. For an exothermic forward reaction, raising temperature decreases K (shifts equilibrium toward reactants). For an endothermic forward reaction, raising temperature increases K.
You cannot change K by changing concentrations, adding a catalyst, or changing pressure alone. Those shift the equilibrium position (which concentrations change) but not the value of K itself.
General equilibrium expression
K = \frac{[C]^c [D]^d}{[A]^a [B]^b}Relationship between K and rate constants
K = \frac{k_f}{k_r}Kp for H₂ + I₂ ⇌ 2 HI
K_p = \frac{(P_{HI})^2}{P_{H_2} \cdot P_{I_2}}Arrhenius equation (how rate constants depend on temperature)
k = A \, e^{-E_a / RT}where A is the pre-exponential factor, Ea is activation energy, R is the gas constant, and T is temperature in Kelvin.
The Haber process for manufacturing ammonia (N₂ + 3 H₂ ⇌ 2 NH₃) is managed entirely around equilibrium principles. Engineers choose high pressure and moderate temperature to push K and the equilibrium position toward products while keeping the reaction rate practical.
Haemoglobin binding oxygen in your lungs and releasing it in your tissues is an equilibrium process. The different partial pressures of O₂ in the lungs vs. the muscles shift the equilibrium in opposite directions.
"At equilibrium nothing is happening." Wrong. Both forward and reverse reactions continue. The rates are equal, not zero. This is the single most common misunderstanding.
"K depends on initial concentrations." It does not. Different starting amounts lead to different equilibrium concentrations, but the ratio expressed by K stays the same (at constant temperature). The simulation data confirms this: trials 2 through 5 all yield K values near 3, despite very different starting conditions.
"Adding a catalyst changes K." A catalyst speeds up both forward and reverse reactions equally. It helps the system reach equilibrium faster but does not shift the equilibrium position or change K.
"K = k_f + k_r." Students sometimes add the rate constants. The correct relationship is K = k_f / k_r (a ratio, not a sum).
⚠️ Be able to write the equilibrium expression for any balanced equation and calculate K or Kp from given concentrations or partial pressures.
⚠️ Know that K = k_f / k_r and be able to explain why K is independent of initial concentrations.
⚠️ Understand the difference between "the reaction has stopped" and "the reaction is at equilibrium." Exams frequently test whether you can distinguish a static situation from dynamic equilibrium.
⚠️ Know that only temperature changes K. Concentration changes, pressure changes, and catalysts affect the position of equilibrium or the speed of reaching it, but not the value of K.
⚠️ Be comfortable computing Kp from partial pressures (plug in, square the HI term, divide). Arithmetic errors on powers of 10 are the most common way to lose marks on this type of calculation.
True or False: At equilibrium, the forward reaction has stopped.
False. Both forward and reverse reactions continue at equal rates.
True or False: Changing the initial concentration of a reactant changes the value of K.
False. K depends only on temperature.
Fill in the blank: K = ___ / k_r.
k_f
True or False: A catalyst changes the value of K.
False. A catalyst speeds up both directions equally.
If K is very large (say, 1380), are products or reactants favoured at equilibrium?
Products are heavily favoured.
Q: Explain what "dynamic equilibrium" means and why it is different from a reaction that has simply stopped.
A: Dynamic equilibrium means both the forward and reverse reactions are still occurring, but at equal rates. The net concentrations of reactants and products do not change. A stopped reaction has a rate of zero in both directions; a dynamic equilibrium has equal, non-zero rates.
Q: For the reaction H₂(g) + I₂(g) ⇌ 2 HI(g), write the Kp expression and calculate Kp given P_H₂ = 1.40 × 10⁻² atm, P_I₂ = 3.50 × 10⁻³ atm, P_HI = 2.60 × 10⁻¹ atm.
A: Kp = (P_HI)² / (P_H₂ × P_I₂) = (0.260)² / (0.0140 × 0.00350) = 0.0676 / 4.90 × 10⁻⁵ ≈ 1380.
Q: In the simulation, k_f = 0.3 and k_r = 0.1. What is the theoretical K, and why do the measured K values from different trials not match it exactly?
A: K = k_f / k_r = 0.3 / 0.1 = 3. The measured values (1.36, 2.56, 3.82, 3.33) scatter around 3 because the simulation uses a stochastic (probabilistic) model that introduces random variation, much like real molecular collisions.
Q: In a real laboratory, you cannot set rate constants directly. What single parameter controls them, and how does changing it affect K?
A: Temperature. Rate constants follow the Arrhenius equation (k = Ae^(−Ea/RT)). Raising temperature increases both k_f and k_r, but by different amounts depending on their activation energies, so K changes. For an exothermic forward reaction, raising T decreases K; for an endothermic forward reaction, raising T increases K.
Q: A student claims that starting with higher concentrations of reactants will increase K. Is this correct? Explain.
A: Incorrect. Higher initial reactant concentrations will shift the equilibrium position so that more product forms at equilibrium, but the ratio of concentrations expressed by K remains the same at a given temperature. K is independent of initial concentrations.
This material connects directly to acid-base chemistry, where Ka and Kb are equilibrium constants for weak acid and weak base dissociation. The idea that K depends only on temperature reappears in thermodynamics (ΔG° = −RT ln K), linking equilibrium to free energy. Solubility product (Ksp) is another special case of the equilibrium constant for dissolution reactions. Le Chatelier's principle, while only lightly touched here, is the qualitative tool you will use throughout general chemistry to predict how systems respond to disturbances.
chemical equilibrium, dynamic equilibrium, equilibrium constant, K, Kp, Kc, Keq, rate constant, forward rate, reverse rate, k_f, k_r, Arrhenius equation, temperature dependence of K, Le Chatelier's principle, equilibrium position, equilibrium shift, ICE table, reaction quotient Q, partial pressure, Haber process, concentration vs. time graph, CHEM 106, NYU chemistry, Lab 9, reactants and products simulation