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Quizzes > Science > Chemistry

Take the Chemistry Equilibrium Test

Ready to ace this chemistry equilibrium quiz? Dive in now!

Difficulty: Moderate
2-5mins
Learning OutcomesCheat Sheet
Paper cut test elements with molecular structures flasks scales layered on teal background for chemistry equilibrium quiz

This equilibrium test helps you practice core chemical equilibrium skills like Le Chatelier's principle, K vs Q, Kc and Kp, ICE tables, and predicting concentration shifts. Start with a warm-up quiz for instant feedback, then tackle the full set at your pace.

Which statement best describes chemical equilibrium in a closed system?
The rates of forward and reverse reactions are equal
All reactants are completely consumed
No products remain in the system
The reaction has stopped completely
At chemical equilibrium, the forward and reverse reaction rates are equal, leading to constant concentrations of reactants and products despite ongoing processes. The reaction does not stop but proceeds in both directions at the same rate, so macroscopic properties remain unchanged. This dynamic balance is characteristic of closed systems where no substances are added or removed.
What is the equilibrium constant expression (Kc) for the reaction: N2 + 3H2 ? 2NH3?
[NH3]^2 / ([N2][H2]^3)
[N2][H2]^3 / [NH3]^2
[NH3] / ([N2][H2])
[NH3]^3 / ([N2]^2[H2])
For N2 + 3H2 ? 2NH3, Kc is products over reactants raised to stoichiometric coefficients: [NH3]^2 divided by [N2][H2]^3. This ratio is constant at a given temperature when equilibrium is established. The correct expression follows directly from the balanced equation.
If the reaction quotient Q is greater than the equilibrium constant K, which way does the reaction proceed to reach equilibrium?
Shift toward reactants
Shift toward products
Remain unchanged
Proceed irreversibly
When Q > K, the concentration of products is too high relative to reactants, so the reaction shifts toward reactants to reestablish equilibrium. This reduces product concentration and increases reactant concentration until Q equals K. Le Chatelier's principle governs this shift.
According to Le Chatelier’s principle, what happens when you increase the concentration of a reactant in a system at equilibrium?
Reaction shifts toward products
Reaction shifts toward reactants
Equilibrium constant changes
Temperature decreases
Le Chatelier’s principle states that when a system at equilibrium is disturbed, it shifts to counteract the disturbance. Adding more reactant causes the system to consume the added reactant by shifting toward products. This shift restores equilibrium ratios without altering the equilibrium constant.
What is the reaction quotient Q for a reaction A + B ? C + D in terms of molar concentrations?
[C][D] / ([A][B])
[A][B] / ([C][D])
[C] + [D] - ([A] + [B])
[A] + [B] / ([C] + [D])
The reaction quotient Q has the same form as the equilibrium constant expression but uses current concentrations. For A + B ? C + D, Q = [C][D] / ([A][B]). It indicates the system’s position relative to equilibrium.
Which equilibrium is an example of a homogeneous equilibrium?
H2(g) + I2(g) ? 2HI(g)
CaCO3(s) ? CaO(s) + CO2(g)
AgCl(s) ? Ag+(aq) + Cl–(aq)
NH4Cl(s) ? NH3(g) + HCl(g)
A homogeneous equilibrium involves all species in the same phase. In H2(g) + I2(g) ? 2HI(g), reactants and products are all gases. Heterogeneous examples have solids or liquids present.
What is the effect on equilibrium when the concentration of a product is decreased?
Shifts toward products
Shifts toward reactants
No change
Equilibrium constant changes
Removing product disturbs equilibrium, so the system shifts toward products to replace what was removed, per Le Chatelier’s principle. This increases product concentration until equilibrium is restored. The equilibrium constant remains unchanged.
What does it mean when a system is in dynamic equilibrium?
Forward and reverse reactions continue at equal rates
Reaction has ceased
Products are completely formed
Reactants are in excess
Dynamic equilibrium means that both forward and reverse reactions occur continuously but at the same rate, so concentrations stay constant. The term dynamic highlights ongoing molecular activity. Macroscopic properties remain unchanged.
How is Kp related to Kc for a gaseous reaction?
Kp = Kc(RT)^(?n)
Kp = Kc / (RT)^(?n)
Kp = Kc RT
Kp = Kc^(RT)
For gas-phase reactions, Kp and Kc are related by Kp = Kc(RT)^(?n), where ?n is moles of gaseous products minus reactants. R is the gas constant and T is temperature in Kelvin. This relation accounts for pressure-concentration conversion.
For the reaction NO2 ? N2O4, Kc at 25°C is 6.8. What is the concentration of NO2 at equilibrium in a 1.0 L flask starting with 1.0 M NO2?
0.63 M
0.85 M
0.42 M
1.37 M
Let x be N2O4 formed: equilibrium [NO2] = 1–2x and [N2O4] = x. Kc = x/(1–2x)^2 = 6.8. Solving gives x ?0.185, so [NO2]?0.63 M. This uses an ICE table and quadratic solution.
What disturbance occurs if temperature is increased for an endothermic equilibrium?
Shifts toward products
Shifts toward reactants
K decreases
Pressure increases
Adding heat to an endothermic reaction is like adding reactant energy, so the equilibrium shifts toward products to absorb excess heat. This raises the equilibrium constant K for the endothermic process. Temperature is the only factor that changes K.
What happens to the equilibrium of a gaseous system when volume is decreased?
Shifts toward side with fewer moles of gas
Shifts toward side with more moles of gas
No shift occurs
K changes sign
Reducing volume increases pressure, so the system shifts toward the side with fewer gaseous moles to counteract the increase. The equilibrium constant remains unchanged because temperature is constant. This response follows Le Chatelier’s principle.
Which expression gives ?G° in terms of the equilibrium constant K?
?G° = –RT ln K
?G° = RT ln K
?G° = –K/RT
?G° = KRT
The relationship ?G° = –RT ln K links thermodynamics with equilibrium. R is gas constant and T is temperature in Kelvin. A larger K gives a more negative ?G°, indicating spontaneity.
In an ICE table, what does the 'C' stand for?
Change in concentration
Concentration at equilibrium
Coefficient in reaction
Constant Kc
ICE stands for Initial, Change, and Equilibrium concentrations. 'Change' represents how much reactant or product concentration changes to reach equilibrium. It is used to calculate equilibrium values via algebra.
What is the reaction quotient Qp for the gas reaction 2SO2 + O2 ? 2SO3?
P_SO3^2 / (P_SO2^2 P_O2)
P_SO2^2 P_O2 / P_SO3^2
P_SO3 / (P_SO2 P_O2)
P_SO2 P_O2 / P_SO3
For gases, Qp uses partial pressures. The expression matches stoichiometry: products over reactants with exponents. Here Qp = (P_SO3)^2 / ((P_SO2)^2 * P_O2). Qp predicts the shift relative to Kp.
If ?G for a reaction is positive, what can be said about Q relative to K?
Q < K
Q > K
Q = K
Q and K are unrelated
When ?G > 0, the reaction is nonspontaneous as written. That means Q < K, so the reaction must proceed forward to reach equilibrium. Conversely, if Q > K, ?G would be negative.
A 0.10 M solution of AgNO3 is mixed with 0.10 M NaCl. Silver chloride (Ksp = 1.8×10^-10) may precipitate. Which compound precipitates first upon adding Cl??
AgCl
NaCl
AgNO3
No precipitation
AgCl has a very low Ksp and will exceed its solubility product before other salts. Adding Cl? to Ag? solution forms AgCl precipitate first once the product [Ag+][Cl–] exceeds 1.8×10^-10. NaCl and AgNO3 are highly soluble.
Which change lowers the value of the equilibrium constant K for an endothermic reaction?
Decreasing temperature
Adding a catalyst
Removing inert gas
Changing pressure at constant temperature
For endothermic reactions, increasing temperature raises K, so decreasing temperature lowers K. Catalysts and pressure changes at constant temperature do not change K. Only temperature affects equilibrium constants.
What effect does a catalyst have on the equilibrium position of a reaction?
No change in position, only rate increases
Shifts toward reactants
Shifts toward products
Changes equilibrium constant
A catalyst lowers activation energies of forward and reverse reactions equally, speeding attainment of equilibrium without changing its position or K. It does not affect reactant or product concentrations at equilibrium.
When an inert gas is added at constant volume to a gas-phase equilibrium, what happens?
No shift occurs
Shifts toward products
Shifts toward reactants
K changes
Adding an inert gas at constant volume raises total pressure but does not change partial pressures of reactants or products. Because the reaction quotient is unchanged, there is no shift and K remains constant.
Which statement about activities in equilibrium expressions is correct?
Activities correct for non-ideal behavior
Activities equal concentrations always
Activities are unitless and ignored
Activities only apply to gases
Activity represents an effective concentration accounting for non-ideal interactions in solution or gas. It differs from molar concentration or pressure especially at high ionic strength. Using activities makes equilibrium constants truly constant.
A 0.05 M solution of a weak acid HA (Ka = 1.0×10^-5) dissociates. What is the pH of the solution?
2.62
3.00
4.00
1.70
Use x^2/(0.05–x)=1.0×10^-5. Assuming x ? 0.05 gives x??(5×10^-7)=7.07×10^-4; pH=–log(7.07×10^-4)=2.65 (approx 2.62). This weak acid approximation is common.
Which graphical plot shows the point of equilibrium on a Gibbs free energy vs. reaction coordinate diagram?
Lowest Gibbs energy point
Highest Gibbs energy point
Where ?G is most positive
Where ?G equals activation energy
Equilibrium corresponds to the reaction coordinate where the Gibbs free energy curve is at its minimum. At that point, ?G for the reaction is zero and the system is most stable. Reactants and products coexist at this energy minimum.
How does the Van ’t Hoff equation relate changes in the equilibrium constant to temperature?
ln(K2/K1) = –?H°/R (1/T2 – 1/T1)
K2/K1 = ?H°/R (T2 – T1)
?K = –R ln(T2/T1)
ln(K2) = ?G°/R T
The Van ’t Hoff equation ln(K2/K1) = –?H°/R (1/T2 – 1/T1) quantifies how K changes with temperature for a reaction of known enthalpy change ?H°. It assumes ?H° is constant over the temperature range. This allows prediction of equilibrium shifts.
When using fugacity instead of pressure in a gas-phase equilibrium, what property does fugacity correct for?
Non-ideal gas behavior
Temperature dependence
Catalyst presence
Volume changes
Fugacity accounts for deviations from ideal gas behavior by providing an effective pressure. It replaces partial pressures in equilibrium expressions at high pressures or non-ideal conditions. Fugacity approaches pressure at low densities.
In a buffer solution, the Henderson-Hasselbalch equation relates pH to pKa and ratio of conjugate base to acid. Which form is correct?
pH = pKa + log([A–]/[HA])
pH = pKa – log([A–]/[HA])
pH = pKa + [A–]/[HA]
pH = pKa – [HA]/[A–]
The Henderson-Hasselbalch equation pH = pKa + log([A–]/[HA]) links the pH of a buffer to the pKa of the acid and the ratio of base to acid. It derives from the acid dissociation constant definition. It predicts pH changes upon mixing acids and bases.
Which activity coefficient model accounts for ionic strength in electrolyte solutions?
Debye–Hückel equation
Ideal Gas Law
Henry’s Law
Arrhenius equation
The Debye–Hückel equation estimates activity coefficients of ions as a function of ionic strength and charge. It corrects for electrostatic interactions in dilute electrolyte solutions. It is critical for accurate equilibrium calculations in non-ideal solutions.
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Study Outcomes

  1. Understand Chemical Equilibrium Principles -

    Grasp the fundamental concepts behind equilibrium in chemical reactions, including dynamic balance and reaction quotients in this equilibrium test.

  2. Apply Le Chatelier's Principle -

    Determine how changes in concentration, pressure, or temperature shift equilibrium positions using this chemistry equilibrium quiz.

  3. Calculate Equilibrium Constants -

    Perform calculations for Kc and Kp values based on given concentration and pressure data in the chemical equilibrium quiz.

  4. Analyze Reaction Dynamics -

    Evaluate how various stress factors influence reaction rates and equilibrium states in an engaging equilibrium quiz format.

  5. Interpret Quiz Feedback -

    Review scored results to identify strengths and weaknesses, guiding further study of chemical equilibrium concepts.

  6. Predict Equilibrium Shifts -

    Use the principles learned to forecast the direction of equilibrium shifts under different experimental conditions.

Cheat Sheet

  1. Le Chatelier's Principle -

    Understanding how stress (concentration, temperature, pressure) shifts equilibrium is key in any equilibrium test. According to IUPAC definitions and university courses, increasing concentration of a reactant pushes the reaction toward products, while raising temperature favors the endothermic direction. A helpful mnemonic: "CAT - Concentration Add, Temperature Endothermic" to remember temperature effects.

  2. Equilibrium Constant (Kc and Kp) -

    The equilibrium constant expresses the ratio of product concentrations to reactant concentrations at equilibrium, and you'll see it in every chemistry equilibrium quiz. For gases, Kp uses partial pressures (Kp = Kc(RT)Δn), while Kc is based on molar concentrations. Remember that a K much greater than 1 means products are favored, whereas K much less than 1 favors reactants.

  3. Reaction Quotient (Q) vs. K -

    Before diving into the chemical equilibrium quiz, always calculate Q to predict direction: Q = ([C]^c[D]^d)/([A]^a[B]^b). If Q < K, the system shifts right (toward products); if Q > K, it shifts left (toward reactants). This quick comparison is indispensable for troubleshooting equilibrium problems on your test.

  4. ICE Tables for Solving Equilibria -

    ICE (Initial, Change, Equilibrium) tables provide a structured way to track concentration changes and solve for unknowns in an equilibrium quiz. List initial concentrations, express changes using ±x, and set up the K expression to solve the resulting equation. Purdue University's chemistry department offers great practice problems to master this technique.

  5. Common Ion Effect & Solubility Equilibria -

    Adding a common ion shifts solubility equilibria and reduces solubility - an important concept in both equilibrium test questions and real-world applications like water treatment. The solubility product constant (Ksp) predicts when a precipitate forms: if [ion] products exceed Ksp, precipitation occurs. Remember "Less Soluble with More Ions" to ace related quiz items.

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