Chemical equilibrium
61 flashcards covering Chemical equilibrium for the MCAT Chem / Phys / Psych / Soc section.
Chemical equilibrium is a fundamental concept in chemistry where a reversible reaction reaches a stable state. At this point, the rate of the forward reaction equals the rate of the reverse reaction, meaning the concentrations of reactants and products remain constant over time. This doesn't mean the reaction has stopped; instead, it's dynamic, with molecules continuing to convert back and forth. Understanding equilibrium helps explain phenomena like acid-base reactions, solubility, and even processes in living organisms, making it essential for grasping how chemical systems behave in real-world scenarios.
On the MCAT, chemical equilibrium appears in the Chem/Phys section, often through multiple-choice questions that test your ability to calculate equilibrium constants (like Kc or Kp), apply Le Chatelier's principle to predict shifts due to changes in concentration, temperature, or pressure, and interpret graphs of reaction progress. Common traps include confusing equilibrium with completion of a reaction or overlooking how catalysts affect rates without shifting the equilibrium position. Focus on mastering calculations, qualitative predictions, and integrating this with biological contexts, such as enzyme kinetics or pH balance.
Practice drawing reaction diagrams to visualize shifts in equilibrium.
Terms (61)
- 01
Chemical Equilibrium
Chemical equilibrium is a state in a reversible reaction where the rates of the forward and reverse reactions are equal, resulting in no net change in the concentrations of reactants and products.
- 02
Reversible Reaction
A reversible reaction is one that can proceed in both the forward and reverse directions, allowing the system to reach a state of equilibrium under the right conditions.
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Dynamic Equilibrium
Dynamic equilibrium occurs when a reversible reaction is ongoing at the molecular level, but the concentrations of reactants and products remain constant over time.
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Law of Mass Action
The law of mass action states that the rate of a chemical reaction is proportional to the product of the concentrations of the reactants, each raised to the power of their stoichiometric coefficients.
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Equilibrium Constant (Kc)
Kc is the equilibrium constant expressed in terms of concentrations, calculated as the ratio of the product of product concentrations to reactant concentrations, each raised to their stoichiometric coefficients, at equilibrium.
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Equilibrium Constant (Kp)
Kp is the equilibrium constant for gas-phase reactions, expressed in terms of partial pressures of the gases involved, similar to Kc but using pressures instead of concentrations.
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Relationship Between Kc and Kp
Kc and Kp are related by the equation Kp = Kc (RT)^Δn, where R is the gas constant, T is the temperature in Kelvin, and Δn is the change in moles of gas from reactants to products.
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Reaction Quotient (Q)
Q is a value calculated similarly to the equilibrium constant but using current concentrations or pressures, which helps predict whether a reaction will proceed forward, reverse, or is at equilibrium.
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Le Chatelier's Principle
Le Chatelier's principle states that if a system at equilibrium is subjected to a change in concentration, pressure, or temperature, the equilibrium will shift to counteract that change.
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Effect of Concentration on Equilibrium
Changing the concentration of a reactant or product causes the equilibrium to shift: increasing a reactant's concentration shifts it toward products, while increasing a product's concentration shifts it toward reactants.
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Effect of Pressure on Equilibrium
For reactions involving gases, increasing pressure shifts the equilibrium toward the side with fewer moles of gas, while decreasing pressure shifts it toward the side with more moles of gas.
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Effect of Temperature on Equilibrium
Temperature changes affect equilibrium based on whether the reaction is exothermic or endothermic: increasing temperature shifts an exothermic reaction toward reactants and an endothermic one toward products.
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Catalysts and Equilibrium
Catalysts speed up both forward and reverse reactions equally, allowing equilibrium to be reached faster, but they do not change the position of the equilibrium or the value of the equilibrium constant.
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ICE Table
An ICE table is a systematic method to organize and solve equilibrium problems, with rows for initial concentrations, changes, and equilibrium concentrations of reactants and products.
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Solubility Product Constant (Ksp)
Ksp is the equilibrium constant for the dissolution of a sparingly soluble ionic compound, representing the product of the concentrations of its ions, each raised to the power of their stoichiometric coefficients.
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Common Ion Effect
The common ion effect occurs when the solubility of a salt is reduced by the presence of a common ion from another source, shifting the equilibrium toward the solid form.
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Saturated Solution
A saturated solution is one in equilibrium with undissolved solute, meaning it contains the maximum amount of solute that can dissolve at that temperature.
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Acid Dissociation Constant (Ka)
Ka is the equilibrium constant for the dissociation of a weak acid in water, equal to the product of the concentrations of the hydrogen ions and the conjugate base divided by the concentration of the undissociated acid.
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Base Dissociation Constant (Kb)
Kb is the equilibrium constant for the dissociation of a weak base in water, equal to the product of the concentrations of the hydroxide ions and the conjugate acid divided by the concentration of the undissociated base.
- 20
Henderson-Hasselbalch Equation
The Henderson-Hasselbalch equation calculates the pH of a buffer solution as pH = pKa + log([conjugate base]/[weak acid]), where pKa is the negative log of Ka.
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Buffer Solutions
Buffer solutions resist changes in pH when small amounts of acid or base are added, typically consisting of a weak acid and its conjugate base or a weak base and its conjugate acid.
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Buffer Capacity
Buffer capacity is the measure of a buffer's ability to resist pH changes, depending on the concentrations of the weak acid and its conjugate base, with higher concentrations providing greater capacity.
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pKa and pKb
pKa is the negative logarithm of Ka for an acid, and pKb is the negative logarithm of Kb for a base, used to indicate the strength of acids and bases respectively.
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Weak Acid Equilibrium
In weak acid equilibrium, only a small fraction of the acid dissociates in water, and the equilibrium is governed by Ka, allowing calculation of pH using the acid's initial concentration.
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Weak Base Equilibrium
Weak base equilibrium involves partial dissociation in water, governed by Kb, which is used to calculate the pH of the solution based on the base's initial concentration.
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Autoionization of Water
Autoionization of water is the equilibrium process where water molecules react to form hydronium and hydroxide ions, with the equilibrium constant Kw equal to 1.0 × 10^-14 at 25°C.
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Polyprotic Acids
Polyprotic acids can donate more than one proton, having multiple dissociation steps with successive Ka values, and their equilibria must be considered stepwise for accurate pH calculations.
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Amphoteric Substances
Amphoteric substances can act as both acids and bases, donating or accepting protons depending on the pH of the solution, such as water or bicarbonate ions.
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Salt Hydrolysis
Salt hydrolysis is the reaction of a salt's ions with water, affecting the pH of the solution, such as when the cation of a weak base hydrolyzes to produce an acidic solution.
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Titration and Equivalence Point
In titration, the equivalence point is reached when the moles of titrant added equal the moles of substance being titrated, often determined by a pH change at equilibrium.
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Indicators in Acid-Base Titrations
Acid-base indicators are weak acids that change color over a specific pH range, used to visually detect the equivalence point in titrations by shifting their equilibrium form.
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Equivalence Point vs. Endpoint
The equivalence point is the theoretical point in a titration where reactants are stoichiometrically equal, while the endpoint is the observed point where the indicator changes color, ideally close to it.
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Solubility and Temperature
Solubility of most solids increases with temperature, affecting the equilibrium of dissolution reactions, which must be considered when predicting precipitation or dissolution.
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Common Ion Effect on Solubility
The common ion effect reduces the solubility of a salt by shifting its dissolution equilibrium toward the solid when a common ion is present, as per Le Chatelier's principle.
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Strategy for Solving Equilibrium Problems
To solve equilibrium problems, first write the balanced equation, set up an ICE table, express the equilibrium constant expression, and solve for unknowns using algebra.
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Calculating K from Initial Concentrations
To calculate K from initial concentrations, use an ICE table to determine equilibrium concentrations after assuming a change and solving the resulting equation based on the reaction's stoichiometry.
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Shifting Equilibrium
Shifting equilibrium refers to how changes in conditions cause the system to move toward a new equilibrium position, as predicted by Le Chatelier's principle.
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Endothermic vs. Exothermic Reactions and Temperature
For endothermic reactions, increasing temperature favors the forward reaction, increasing K, while for exothermic reactions, it favors the reverse, decreasing K.
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Partial Pressures in Kp
In Kp expressions, partial pressures of gases are used instead of concentrations, calculated as the mole fraction of the gas times the total pressure.
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Heterogeneous Equilibria
Heterogeneous equilibria involve reactants and products in different phases, such as a solid and a gas, where pure solids and liquids are not included in the equilibrium constant expression.
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Homogeneous Equilibria
Homogeneous equilibria occur when all reactants and products are in the same phase, such as all gases or all in solution, simplifying the equilibrium constant expression.
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Gibbs Free Energy and Equilibrium
The Gibbs free energy change, ΔG, relates to the equilibrium constant by ΔG = -RT ln K, where a negative ΔG indicates a reaction that proceeds spontaneously toward equilibrium.
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Relationship Between K and ΔG°
The standard Gibbs free energy change, ΔG°, is related to K by ΔG° = -RT ln K, allowing prediction of equilibrium constants from thermodynamic data.
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Q vs. K: Predicting Reaction Direction
If Q is less than K, the reaction proceeds forward; if Q is greater than K, it proceeds reverse; if Q equals K, the system is at equilibrium.
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Equilibrium in Gas Reactions
In gas reactions, equilibrium is influenced by partial pressures and can be expressed with Kp, requiring consideration of the ideal gas law for accurate calculations.
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Equilibrium in Solution
Equilibrium in solution typically involves aqueous reactions, where concentrations are used for Kc, and factors like ionic strength can affect the position of equilibrium.
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Common Mistakes with Le Chatelier
A common mistake is assuming catalysts shift equilibrium, but they only affect the rate; another is ignoring that volume changes only impact gas equilibria.
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Solubility Rules
Solubility rules are guidelines to predict whether ionic compounds will dissolve in water, based on empirical observations, helping determine if a precipitation reaction will occur at equilibrium.
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Precipitation Reactions
Precipitation reactions occur when two aqueous solutions form an insoluble solid, driven by the solubility equilibrium shifting to form the precipitate.
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Complex Ion Formation
Complex ion formation involves a metal ion combining with ligands to form a stable complex, which can shift equilibria like dissolution by removing free metal ions.
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Determining Ksp from Solubility
To determine Ksp from solubility, multiply the ion concentrations at equilibrium, derived from the solubility value, according to the compound's dissociation equation.
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Using Q to Predict Precipitation
Compare Q to Ksp: if Q exceeds Ksp, precipitation occurs; if Q is less than Ksp, no precipitate forms; if equal, the solution is at saturation equilibrium.
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Le Chatelier in Industrial Processes
In industrial processes like the Haber process, Le Chatelier's principle guides conditions such as high pressure and moderate temperature to maximize ammonia yield at equilibrium.
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Temperature Dependence of K
The value of K changes with temperature according to the Van't Hoff equation, increasing for endothermic reactions and decreasing for exothermic ones as temperature rises.
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Pressure Changes in Gas Equilibria
Pressure changes in gas equilibria only affect systems where the number of moles of gas differs between sides, shifting to minimize the pressure change.
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Concentration Changes and Q
Adding or removing a reactant or product changes Q, causing the reaction to shift until Q equals K again, restoring equilibrium.
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Example: Haber Process
In the Haber process, N2 + 3H2 ⇌ 2NH3, increasing pressure shifts equilibrium toward NH3 due to fewer moles of gas on the product side.
At 200 atm and 400°C, the yield of ammonia increases compared to lower pressures.
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Example: Dissociation of Acetic Acid
For acetic acid dissociation, CH3COOH ⇌ H+ + CH3COO-, the equilibrium constant Ka is 1.8 × 10^-5, indicating it's a weak acid with low dissociation.
In a 0.1 M solution, only about 1.3% dissociates at equilibrium.
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Example: Calculating pH of a Buffer
For a buffer of 0.1 M acetic acid and 0.1 M sodium acetate, use Henderson-Hasselbalch: pH = 4.74 + log(1), resulting in pH 4.74.
Adding a small amount of base changes pH minimally due to the buffer.
- 60
Example: AgCl Solubility
For AgCl, Ksp = 1.8 × 10^-10, so its solubility is low, and adding NaCl (common ion) further reduces it by shifting the dissolution equilibrium.
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Example: Q for a Reaction
For 2A + B ⇌ C, if [A] = 0.2 M, [B] = 0.1 M, [C] = 0.05 M, and Kc = 10, then Q = (0.05) / (0.2)^2 (0.1) = 1.25, so reaction proceeds forward.