13.4 Equilibrium Calculations
3 min read•june 25, 2024
is a balancing act in . When forward and reverse rates equalize, concentrations stabilize. This dynamic state can be influenced by external factors, shifting the balance to counteract changes.
constants quantify the relative amounts of reactants and products at equilibrium. ICE tables help organize and solve equilibrium problems, making it easier to calculate concentrations and predict reaction directions.
Equilibrium Concepts
Changes in equilibrium concentrations
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- states that a system at equilibrium will respond to a stress by shifting to counteract the stress and re-establish equilibrium
- Stresses include changes in concentration (adding or removing reactants or products), pressure (increasing or decreasing), volume (increasing or decreasing), or temperature (increasing or decreasing)
- Concentration changes shift equilibrium
- Adding reactants (NaOH) or removing products (PbS) shifts equilibrium to the right (towards products)
- Adding products (NH₃) or removing reactants (N₂) shifts equilibrium to the left (towards reactants)
- Pressure and volume changes affect gaseous equilibria
- Increasing pressure (decreasing volume) shifts equilibrium towards the side with fewer moles of gas (2 NO₂)
- Decreasing pressure (increasing volume) shifts equilibrium towards the side with more moles of gas (N₂ + 3 H₂)
- Temperature changes shift equilibrium in the or direction
- Increasing temperature shifts equilibrium in the endothermic direction (heat is absorbed, ΔH > 0)
- Decreasing temperature shifts equilibrium in the exothermic direction (heat is released, ΔH < 0)
- Catalysts (enzymes) do not affect the position of equilibrium, only increase the rate at which it is reached
- The refers to the relative amounts of reactants and products present at equilibrium
Chemical Equilibrium and Reversible Reactions
- Chemical equilibrium occurs in reversible reactions when the forward and reverse reaction rates become equal
- Reversible reactions can proceed in both directions and are represented by a double arrow (⇌)
- At equilibrium, the system reaches a state of where the concentrations of reactants and products remain constant over time
Calculation of equilibrium constants
- represents the ratio of product concentrations to reactant concentrations at equilibrium
- For a general reaction ,
- Concentrations are raised to the power of their stoichiometric coefficients (a, b, c, d)
- Pure solids (CaCO₃) and liquids (H₂O) are not included in the equilibrium constant expression
- Solving for equilibrium concentrations involves writing the balanced chemical equation, equilibrium constant expression, substituting known concentrations and solving for the unknown concentration
- Relationship between and determines the direction of the reaction
- If , the reaction will proceed to the right (towards products) to reach equilibrium
- If , the reaction will proceed to the left (towards reactants) to reach equilibrium
- If , the system is at equilibrium and no net change occurs
- The can influence equilibrium by shifting the position when an ion already present in the system is added
ICE table for equilibrium problems
- organizes information about concentrations in a reaction at different stages: Initial (before reaction starts), Change (as reaction proceeds), and Equilibrium (final concentrations)
- Steps for using an ICE table:
- Write the balanced chemical equation ()
- Create a table with rows for Initial, Change, and Equilibrium concentrations
- Fill in known initial concentrations ()
- Represent changes in concentration using variables (let change in )
- Write equilibrium concentrations in terms of initial concentrations and change variables ()
- Substitute equilibrium concentrations into the equilibrium constant expression ()
- Solve the resulting equation for the change variable ()
- Calculate equilibrium concentrations using the solved change variable ()
Key Terms to Review (26)
Aqueous solution: An aqueous solution is a solution in which water is the solvent. Commonly used in chemistry to describe reactions occurring in water.
Chemical Equilibrium: Chemical equilibrium is a state where the forward and reverse reactions in a chemical system occur at equal rates, resulting in no net change in the concentrations of the reactants and products over time. This dynamic balance is a fundamental concept in understanding the behavior of chemical systems.
Common Ion Effect: The common ion effect is a principle in chemistry that describes the influence of a common ion on the solubility of a salt or the position of a chemical equilibrium. It is a fundamental concept that underlies various equilibrium processes in chemistry, including equilibrium calculations, hydrolysis of salts, buffer solutions, precipitation and dissolution, and coupled equilibria.
Dynamic equilibrium: Dynamic equilibrium occurs when the rates of the forward and reverse processes are equal, resulting in no net change in the system. It is a key concept in phase transitions where phases coexist at equilibrium.
Dynamic Equilibrium: Dynamic equilibrium is a state in which opposing chemical processes occur at equal rates, resulting in a stable and unchanging overall system. It is a fundamental concept in chemistry that describes the balance between forward and reverse reactions in a closed system.
Endothermic: Endothermic refers to a process or reaction that absorbs heat from the surrounding environment. This means that the system undergoing the endothermic process requires an input of energy in the form of heat in order to proceed. Endothermic processes are central to understanding various topics in chemistry, including energy basics, enthalpy, dissolution, equilibrium, and free energy.
Endothermic process: An endothermic process is a chemical reaction or physical change that absorbs heat energy from its surroundings. These processes result in a decrease in the temperature of the surrounding environment.
Equilibrium: Equilibrium is a state of balance or stability in a system, where the opposing forces or processes are in a state of dynamic balance. It is a fundamental concept that underpins various aspects of chemistry, including phase changes, chemical reactions, and thermodynamic processes.
Equilibrium Constant: The equilibrium constant is a quantitative measure of the extent of a chemical reaction at equilibrium. It represents the ratio of the concentrations of the products to the reactants, raised to their respective stoichiometric coefficients, and is a fundamental concept in understanding the behavior of chemical systems at equilibrium.
Equilibrium constant, K: The equilibrium constant, $K$, is a ratio that quantifies the concentrations of reactants and products in a chemical reaction at equilibrium. It provides insight into the position of the equilibrium and the extent to which reactants are converted into products.
Equilibrium Position: Equilibrium position refers to the state in which the forward and reverse reactions in a chemical system occur at equal rates, resulting in a constant composition of the reactants and products. This term is central to understanding the behavior and characteristics of chemical equilibria.
Exothermic: Exothermic refers to a chemical reaction or process that releases energy in the form of heat to the surrounding environment. These reactions produce more energy than they consume, resulting in a net release of heat.
Exothermic process: An exothermic process is a chemical reaction or physical change that releases heat to its surroundings. This release of energy usually results in an increase in the temperature of the surroundings.
Henry Louis Le Chatelier: Henry Louis Le Chatelier was a French chemist who formulated the principle of chemical equilibrium, known as Le Chatelier's principle. This principle describes how a chemical system at equilibrium responds to changes in the conditions of the system, such as concentration, temperature, or pressure.
Heterogeneous equilibrium: Heterogeneous equilibrium involves reactants and products in different phases, such as solids, liquids, and gases. The equilibrium constant expression for these systems only includes the concentrations of the gaseous and aqueous species.
Heterogeneous Equilibrium: A heterogeneous equilibrium is a state of balance that exists between chemical species in a system where at least one of the reactants or products is in a different physical state, such as a solid, liquid, or gas, from the others. This type of equilibrium is an important concept in the context of chemical equilibria and equilibrium calculations.
Homogeneous equilibrium: Homogeneous equilibrium occurs when all reactants and products of a chemical reaction are in the same phase, usually liquid or gas. It is characterized by a constant ratio of the concentrations of reactants and products at equilibrium.
Homogeneous Equilibrium: A homogeneous equilibrium is a state of balance in a chemical system where the concentrations of all reactants and products remain constant over time, and the system is composed of a single phase, such as a gas or a solution. This term is crucial in understanding the principles of chemical equilibria, equilibrium constants, and equilibrium calculations.
ICE Table: An ICE table, also known as an Initial, Change, and Equilibrium table, is a tool used in chemistry to organize and analyze the concentrations of reactants and products in a chemical equilibrium system. It provides a structured way to visualize and calculate the changes in concentrations as a reaction approaches equilibrium.
Kc: Kc, or the equilibrium constant, is a quantitative measure of the extent of a chemical reaction at equilibrium. It represents the ratio of the concentrations of the products to the concentrations of the reactants, raised to their respective stoichiometric coefficients. The value of Kc provides insight into the position of the equilibrium and the relative amounts of products and reactants present at equilibrium.
Kp: Kp, also known as the equilibrium constant for partial pressures, is a measure of the relative concentrations of reactants and products at equilibrium in a chemical reaction. It is a fundamental concept in the study of chemical equilibria and is used to predict the direction and extent of a reaction under specific conditions.
Le Chatelier's Principle: Le Chatelier's Principle states that when a system at equilibrium is subjected to a change in one of the conditions (concentration, temperature, or pressure) affecting that equilibrium, the system will shift to counteract the change and re-establish equilibrium.
Q: Q, also known as the equilibrium constant, is a measure of the relative concentrations of reactants and products at equilibrium in a chemical reaction. It is a fundamental concept in understanding the behavior of chemical systems and is essential in the study of equilibrium constants, equilibrium calculations, and precipitation and dissolution processes.
Reaction Quotient: The reaction quotient, denoted as Q, is a measure of the relative concentrations of the products and reactants in a chemical reaction at any given time, regardless of whether the system has reached equilibrium or not. It is a useful tool for understanding the direction and extent of a reaction as it progresses towards equilibrium.
Reaction quotient (Q): The reaction quotient, Q, is a measure of the relative amounts of products and reactants present in a reaction mixture at any given point in time. It is calculated using the same expression as the equilibrium constant but with current concentrations or partial pressures.
Reversible Reactions: Reversible reactions are chemical reactions where the forward and backward reactions can occur simultaneously. The products of the forward reaction can react to form the original reactants, and vice versa, until a state of equilibrium is reached.