15: Chemical Equilibrium
- Page ID
- 47429
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)In previous science classes, you may have learned that one way to distinguish chemical changes from physical changes is that physical changes—such as the melting and freezing of water—are reversible, but that chemical changes are not. In this chapter, we will see that this simple answer is not necessarily what it seems.
- 15.2: The Rate of a Chemical Reaction
- This page covers collision theory, factors influencing reaction rates, and reversible reactions. It emphasizes the necessity of collisions with adequate energy and orientation, impacted by temperature, concentration, and surface area. Higher temperatures and concentrations enhance reaction rates, with practical examples like food preservation and fire-starting. Catalysts reduce activation energy without being consumed.
- 15.3: The Idea of Dynamic Chemical Equilibrium
- This page introduces chemistry and its relevance, detailing the scientific method and foundational concepts for beginners. It covers reversible reactions and equilibrium, explaining outcomes when reactants combine and the conditions for reaching dynamic equilibrium.
- 15.4: The Equilibrium Constant - A Measure of How Far a Reaction Goes
- This page explains equilibrium constant expressions in chemical reactions, where concentrations of reactants and products remain constant at equilibrium. It details how to write these expressions, with products in the numerator and reactants in the denominator, using coefficients as exponents. The value of the equilibrium constant \(K\) indicates the balance between reactants and products—greater than 1 favors products, less than 1 favors reactants, and equal to 1 signifies no preference.
- 15.5: Heterogeneous Equilibria- The Equilibrium Expression for Reactions Involving a Solid or a Liquid
- This page discusses building educational maps, emphasizing the time needed to create materials. LibreTexts encourages sharing incomplete pages and invites contributions of open resources to improve the project. It also welcomes reader suggestions through email.
- 15.6: Calculating and Using Equilibrium Constants
- This page explains homogeneous and heterogeneous equilibria in chemical reactions, detailing how equilibrium constants are calculated. It emphasizes that pure solids and liquids have activities treated as 1, simplifying calculations. Examples are provided for writing equilibrium constant expressions, underscoring that solid or liquid phases do not impact equilibrium composition. Additionally, it clarifies that solute activities in solution are determined by their molarities.
- 15.7: Disturbing a Reaction at Equilibrium- Le Châtelier’s Principle
- This page explains Le Chatelier's Principle, which states that an equilibrium system will adjust to counteract external stresses like changes in concentration or temperature. An increase in reactant concentration shifts the equilibrium towards products, while a decrease shifts it towards reactants. This understanding is crucial for predicting the responses of chemical systems to different conditions, ensuring they maintain dynamic equilibrium.
- 15.8: The Effect of a Concentration Change on Equilibrium
- This page examines Le Chatelier's Principle in a chemical equilibrium involving iron ions (Fe³⁺) and thiocyanate ions (SCN⁻), highlighting how the addition or removal of these ions affects the equilibrium position. Adding Fe³⁺ shifts the equilibrium to the right, forming more FeSCN²⁺ and darkening the solution, while adding FeSCN²⁺ shifts it left. Removing SCN⁻ also causes a left shift. Despite these shifts, the equilibrium constant (Keq) remains constant.
- 15.9: The Effect of a Volume Change on Equilibrium
- This page explains how pressure and volume changes influence chemical equilibrium in gas systems. Increasing pressure shifts equilibrium toward the side with fewer moles of gas, while decreasing pressure favors the side with more moles. If both sides have equal moles, equilibrium remains unchanged. Solids and liquids are unaffected. An example demonstrates that increasing pressure in the reaction N2O4(g) <=> 2NO2(g) shifts equilibrium left, favoring the reactants.
- 15.10: The Effect of Temperature Changes on Equilibrium
- This page explains the impact of temperature on equilibrium systems, noting that higher temperatures favor endothermic reactions, shifting equilibrium to absorb heat. It emphasizes that temperature changes alter the equilibrium constant (Keq), unlike concentration, pressure, or volume changes. Favoring the forward reaction raises Keq, while favoring the reverse reaction lowers it.
- 15.11: The Solubility-Product Constant
- This page covers solution-solid equilibria and the solubility product constant (\(K_{sp}\)), highlighting its role as an equilibrium constant for saturated solutions. It explains the equilibrium expression, which omits solids and centers on ion concentrations, and provides examples for deriving \(K_{sp}\) and calculating ion concentrations.
- 15.12: The Path of a Reaction and the Effect of a Catalyst
- This page explains that adding a catalyst to an equilibrium system accelerates both the forward and reverse reactions equally, serving as a stress factor but not changing the equilibrium position or the ratio of reactants to products. While a catalyst speeds up the achievement of equilibrium, it does not alter the overall equilibrium state.