Chemical equilibrium is a dynamic state in which opposing processes occur at the same rate. In this unit, students learn that any bond or intermolecular attraction that can be formed can be broken. These two processes are in a dynamic competition, sensitive to initial conditions and external perturbations. A change in conditions, such as addition of a chemical species, change in temperature, or change in volume, can cause the rate of the forward and reverse reactions to fall out of balance. Le Châtelier’s principle provides a means to reason qualitatively about the direction of the shift in an equilibrium system resulting from various possible stresses. The expression for the equilibrium constant, K, is a mathematical expression that describes the equilibrium state associated with a chemical change. An analogous expression for the reaction quotient, Q, describes a chemical reaction at any point, enabling a comparison to the equilibrium state. Subsequent units will explore equilibrium constants that arise from acid-base chemistry.
3.A Represent chemical phenomena using appropriate graphing techniques, including correct scale and units.
3.C Represent visually the relationship between the structures and interactions across multiple levels or scales (e.g., particulate to macroscopic).
4.D Explain the degree to which a model or representation describes the connection between particulate-level properties and macroscopic properties.
5.A Identify quantities needed to solve a problem from given information (e.g., text, mathematical expressions, graphs, or tables).
5.C Explain the relationship between variables within an equation when one variable changes.
5.F Calculate, estimate, or predict an unknown quantity from known quantities by selecting and following a logical computational pathway and attending to precision (e.g., performing dimensional analysis and attending to significant figures).
6.D Provide reasoning to justify a claim using chemical principles or laws, or using mathematical justification.
6.F Explain the connection between experimental results and chemical concepts, processes, or theories.
Explain the relationship between the occurrence of a reversible chemical or physical process, and the establishment of equilibrium, to experimental observations.
Many observable processes are reversible. Examples include evaporation and condensation of water, absorption and desorption of a gas, or dissolution and precipitation of a salt. Some important reversible chemical processes include the transfer of protons in acid-base reactions and the transfer of electrons in redox reactions.
When equilibrium is reached, no observable changes occur in the system. Reactants and products are simultaneously present, and the concentrations or partial pressures of all species remain constant.
The equilibrium state is dynamic. The forward and reverse processes continue to occur at equal rates, resulting in no net observable change.
Graphs of concentration, partial pressure, or rate of reaction versus time for simple chemical reactions can be used to understand the establishment of chemical equilibrium.
Explain the relationship between the direction in which a reversible reaction proceeds and the relative rates of the forward and reverse reactions.
If the rate of the forward reaction is greater than the reverse reaction, then there is a net conversion of reactants to products. If the rate of the reverse reaction is greater than that of the forward reaction, then there is a net conversion of products to reactants. An equilibrium state is reached when these rates are equal.
Represent the reaction quotient Qc or Qp , for a reversible reaction, and the corresponding equilibrium expressions Kc = Qc or K p = Qp .
The reaction quotient Qc describes the relative concentrations of reaction species at any time. For gas phase reactions, the reaction quotient may instead be written in terms of pressures as Qp . The reaction quotient tends toward the equilibrium constant such that at equilibrium Kc = Qc and K p = Qp .
The reaction quotient does not include substances whose concentrations (or partial pressures) are independent of the amount, such as for solids and pure liquids.
Calculate Kc or K p based on experimental observations of concentrations or pressures at equilibrium.
Equilibrium constants can be determined from experimental measurements of the concentrations or partial pressures of the reactants and products at equilibrium.
Explain the relationship between very large or very small values of K and the relative concentrations of chemical species at equilibrium.
Some equilibrium reactions have very large K values and proceed essentially to completion. Others have very small K values and barely proceed at all.
Represent a multistep process with an overall equilibrium expression, using the constituent K expressions for each individual reaction.
When a reaction is reversed, K is inverted.
When the stoichiometric coefficients of a reaction are multiplied by a factor c, K is raised to the power c
When reactions are added together, the K of the resulting overall reaction is the product of the K’s for the reactions that were summed.
Since the expressions for K and Q have identical mathematical forms, all valid algebraic manipulations of K also apply to Q.
Identify the concentrations or partial pressures of chemical species at equilibrium based on the initial conditions and the equilibrium constant.
The concentrations or partial pressures of species at equilibrium can be predicted given the balanced reaction, initial concentrations, and the appropriate K.
Represent a system undergoing a reversible reaction with a particulate model.
Particulate representations can be used to describe the relative numbers of reactant and product particles present prior to and at equilibrium, and the value of the equilibrium constant.
Identify the response of a system at equilibrium to an external stress, using Le Châtelier's principle.
Le Châtelier’s principle can be used to predict the response of a system to stresses such as addition or removal of a chemical species, change in temperature, change in volume/ pressure of a gas-phase system, or dilution of a reaction system.
Le Châtelier’s principle can be used to predict the effect that a stress will have on experimentally measurable properties such as pH, temperature, and color of a solution.
Explain the relationships between Q, K, and the direction in which a reversible reaction will proceed to reach equilibrium.
A disturbance to a system at equilibrium causes Q to differ from K, thereby taking the system out of equilibrium. The system responds by bringing Q back into agreement with K, thereby establishing a new equilibrium state.
Some stresses, such as changes in concentration, cause a change in Q only. A change in temperature causes a change in K. In either case, the concentrations or partial pressures of species redistribute to bring Q and K back into equality.
Calculate the solubility of a salt based on the value of K sp for the salt.
The dissolution of a salt is a reversible process whose extent can be described by K sp, the solubility-product constant.
The solubility of a substance can be calculated from the K sp for the dissolution process. This relationship can also be used to predict the relative solubility of different substances.
The solubility rules can be quantitatively related to Ksp, in which Ksp values >1 correspond to soluble salts.
Identify the solubility of a salt, and/or the value of K sp for the salt, based on the concentration of a common ion already present in solution.
The solubility of a salt is reduced when it is dissolved into a solution that already contains one of the ions present in the salt. The impact of this “common-ion effect” on solubility can be understood qualitatively using Le Châtelier’s principle or calculated from the K sp for the dissolution process.
Identify the qualitative effect of changes in pH on the solubility of a salt.
The solubility of a salt is pH sensitive when one of the constituent ions is a weak acid or base. These effects can be understood qualitatively using Le Châtelier’s principle.
Explain the relationship between the solubility of a salt and changes in the enthalpy and entropy that occur in the dissolution process.
The free energy change (ΔG°) for dissolution of a substance reflects a number of factors: the breaking of the intermolecular interactions that hold the solid together, the reorganization of the solvent around the dissolved species, and the interaction of the dissolved species with the solvent. It is possible to estimate the sign and relative magnitude of the enthalpic and entropic contributions to each of these factors. However, making predictions for the total change in free energy of dissolution can be challenging due to the cancellations among the free energies associated with the three factors cited.