Search

Text Color

Margin Size

Font Type

Enable Dyslexic Font

6: Equilibrium Chemistry

Last updated

Sep 12, 2021
Save as PDF
- 5.9: Chapter Summary and Key Terms
- 6.1: Reversible Reactions and Chemical Equilibria

$\newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} }$

$\newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}}$

$\newcommand{\id}{\mathrm{id}}$ $\newcommand{\Span}{\mathrm{span}}$

( \newcommand{\kernel}{\mathrm{null}\,}\) $\newcommand{\range}{\mathrm{range}\,}$

$\newcommand{\RealPart}{\mathrm{Re}}$ $\newcommand{\ImaginaryPart}{\mathrm{Im}}$

$\newcommand{\Argument}{\mathrm{Arg}}$ $\newcommand{\norm}[1]{\| #1 \|}$

$\newcommand{\inner}[2]{\langle #1, #2 \rangle}$

$\newcommand{\Span}{\mathrm{span}}$

$\newcommand{\id}{\mathrm{id}}$

$\newcommand{\Span}{\mathrm{span}}$

$\newcommand{\kernel}{\mathrm{null}\,}$

$\newcommand{\range}{\mathrm{range}\,}$

$\newcommand{\RealPart}{\mathrm{Re}}$

$\newcommand{\ImaginaryPart}{\mathrm{Im}}$

$\newcommand{\Argument}{\mathrm{Arg}}$

$\newcommand{\norm}[1]{\| #1 \|}$

$\newcommand{\inner}[2]{\langle #1, #2 \rangle}$

$\newcommand{\Span}{\mathrm{span}}$ $\newcommand{\AA}{\unicode[.8,0]{x212B}}$

$\newcommand{\vectorA}[1]{\vec{#1}} % arrow$

$\newcommand{\vectorAt}[1]{\vec{\text{#1}}} % arrow$

$\newcommand{\vectorB}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} }$

$\newcommand{\vectorC}[1]{\textbf{#1}}$

$\newcommand{\vectorD}[1]{\overrightarrow{#1}}$

$\newcommand{\vectorDt}[1]{\overrightarrow{\text{#1}}}$

$\newcommand{\vectE}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash{\mathbf {#1}}}}$

$\newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} }$

$\newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}}$

$\newcommand{\avec}{\mathbf a}$

$\newcommand{\bvec}{\mathbf b}$

$\newcommand{\cvec}{\mathbf c}$

$\newcommand{\dvec}{\mathbf d}$

$\newcommand{\dtil}{\widetilde{\mathbf d}}$

$\newcommand{\evec}{\mathbf e}$

$\newcommand{\fvec}{\mathbf f}$

$\newcommand{\nvec}{\mathbf n}$

$\newcommand{\pvec}{\mathbf p}$

$\newcommand{\qvec}{\mathbf q}$

$\newcommand{\svec}{\mathbf s}$

$\newcommand{\tvec}{\mathbf t}$

$\newcommand{\uvec}{\mathbf u}$

$\newcommand{\vvec}{\mathbf v}$

$\newcommand{\wvec}{\mathbf w}$

$\newcommand{\xvec}{\mathbf x}$

$\newcommand{\yvec}{\mathbf y}$

$\newcommand{\zvec}{\mathbf z}$

$\newcommand{\rvec}{\mathbf r}$

$\newcommand{\mvec}{\mathbf m}$

$\newcommand{\zerovec}{\mathbf 0}$

$\newcommand{\onevec}{\mathbf 1}$

$\newcommand{\real}{\mathbb R}$

$\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}$

$\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}$

$\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}$

$\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}$

$\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}$

$\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}$

$\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}$

$\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}$

$\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}$

$\newcommand{\laspan}[1]{\text{Span}\{#1\}}$

$\newcommand{\bcal}{\cal B}$

$\newcommand{\ccal}{\cal C}$

$\newcommand{\scal}{\cal S}$

$\newcommand{\wcal}{\cal W}$

$\newcommand{\ecal}{\cal E}$

$\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}$

$\newcommand{\gray}[1]{\color{gray}{#1}}$

$\newcommand{\lgray}[1]{\color{lightgray}{#1}}$

$\newcommand{\rank}{\operatorname{rank}}$

$\newcommand{\row}{\text{Row}}$

$\newcommand{\col}{\text{Col}}$

$\renewcommand{\row}{\text{Row}}$

$\newcommand{\nul}{\text{Nul}}$

$\newcommand{\var}{\text{Var}}$

$\newcommand{\corr}{\text{corr}}$

$\newcommand{\len}[1]{\left|#1\right|}$

$\newcommand{\bbar}{\overline{\bvec}}$

$\newcommand{\bhat}{\widehat{\bvec}}$

$\newcommand{\bperp}{\bvec^\perp}$

$\newcommand{\xhat}{\widehat{\xvec}}$

$\newcommand{\vhat}{\widehat{\vvec}}$

$\newcommand{\uhat}{\widehat{\uvec}}$

$\newcommand{\what}{\widehat{\wvec}}$

$\newcommand{\Sighat}{\widehat{\Sigma}}$

$\newcommand{\lt}{<}$

$\newcommand{\gt}{>}$

$\newcommand{\amp}{&}$

$\definecolor{fillinmathshade}{gray}{0.9}$

Regardless of the problem on which an analytical chemist is working, its solution requires a knowledge of chemistry and the ability to apply that knowledge to solve a problem. For example, an analytical chemist who is studying the effect of pollution on spruce trees needs to know, or know where to find, the chemical differences between p‐hydroxybenzoic acid and p‐hydroxyacetophenone, two common phenols found in the needles of spruce trees.

The ability to “think as a chemist” is a product of your experience in the classroom and in the laboratory. For the most part, the material in this text assumes you are familiar with topics covered in earlier courses; however, because of its importance to analytical chemistry, this chapter provides a review of equilibrium chemistry. Much of the material in this chapter should be familiar to you, although some topics—ladder diagrams and activity, for example—likely afford you with new ways to look at equilibrium chemistry.

6.1: Reversible Reactions and Chemical Equilibria
In 1798, chemist Claude Berthollet observed Na2CO3 deposits at Egypt's Natron Lakes, which contradicted existing chemical theory based on elective affinities that dictated reactions proceed in one direction only. Berthollet's insight into the reversibility of reactions, exemplified by the formation of Na2CO3 using CaCO3 and NaCl, contributed to the understanding of chemical equilibrium.
6.2: Thermodynamics and Equilibrium Chemistry
The page discusses the principles of thermodynamics, focusing on chemical reactions and the factors influencing their equilibrium positions. It outlines the roles of Gibbs free energy, enthalpy, and entropy in determining whether reactions are thermodynamically favorable. The Gibbs free energy equation predicts reaction direction under specific conditions. The equilibrium constant (K) describes a reaction's equilibrium position using concentrations or partial pressures.
6.3: Manipulating Equilibrium Constants
The page elaborates on two principles regarding equilibrium constants: reversing a reaction inverts its equilibrium constant, and combining reactions involves multiplying their equilibrium constants. An example and exercise demonstrate these principles. For the example, the equilibrium constant for a reaction is calculated by combining constants of related reactions, resulting in 0.10. In a similar exercise, the equilibrium constant for a different reaction is calculated to be approximately 31.
6.4: Equilibrium Constants for Chemical Reactions
The document provides an in-depth overview of several essential chemical reactions relevant to analytical chemistry, such as precipitation, acid-base, complexation, and oxidation-reduction (redox) reactions. It explains the concepts of equilibrium constants like Ksp for precipitation reactions, Ka and Kb for acid-base reactions, and Kf for complexation reactions. The text discusses strong and weak acids and bases, amphiprotic species, the dissociation of water, and the pH scale.
6.5: Le Châtelier’s Principle
The document explains the concept of chemical equilibria and Le Ch??telier's principle through examples involving acetic acid dissociation and silver chloride solubility. It discusses how adding reactants or products affects equilibrium, maintaining the equilibrium constant despite changes. It elaborates on how changing the concentration, such as adding sodium acetate or a ligand, affects reactions, and also how pressure and volume changes influence equilibrium through the ideal gas law.
6.6: Ladder Diagrams
The page discusses the importance of considering chemical interactions, like pH and solubility, when developing or evaluating analytical methods. It critiques the inappropriate use of NH3 in precipitating AgCl due to its solubility-increasing effect. Key analytical errors often stem from overlooking chemical interferences. Ladder diagrams are introduced as tools for visualizing equilibrium chemistry, aiding in understanding reaction dynamics and evaluating changes in solution conditions.
6.7: Solving Equilibrium Problems
This page discusses using ladder diagrams and algebraic solutions to evaluate and solve equilibrium problems related to chemical reactivity and solubility. It begins with a straightforward example of calculating the solubility of Pb(IO3)2 in deionized water and proceeds to more complex scenarios considering the common ion effect and the presence of ligands.
6.8: Buffer Solutions
This page explains the different responses to adding HCl to pure water versus a solution with acetic acid and sodium acetate. It describes how buffers, like the acetic acid-sodium acetate mixture, resist changes in pH due to their equilibrium shifting. The Henderson-Hasselbalch equation is central to understanding buffer preparation and effectiveness.
6.9: Activity Effects
The stability of the metal???ligand complex Fe(SCN)2+ decreases in the presence of inert ions, as seen when adding an inert salt like KNO3 to an equilibrium mixture of Fe3+ and SCN???. This leads to a shift in the equilibrium, reducing the concentration of Fe(SCN)2+ and lowering its formation constant, K1. The ionic strength of the solution, which is a measure of the concentration and charge of ions, affects the apparent formation constant.
6.10: Using Excel and R to Solve Equilibrium Problems
The document discusses solving equilibrium problems in chemistry using tools like Excel and R, emphasizing the importance of simplifying assumptions to avoid complex equations. Excel's Solver function helps find polynomial roots and solve simultaneous equations, demonstrated with examples on solubility and pH calculations. R's uniroot command and custom functions can also solve these problems by iterating on solutions.
6.11: Some Final Thoughts on Equilibrium Calculations
The chapter discusses tools for evaluating system composition at equilibrium, highlighting the importance of selecting the appropriate tool based on the precision required. It emphasizes the need to include all relevant equilibrium reactions to prevent errors. It introduces computational programs like Visual Minteq and CurTiPot for modeling equilibria and the R package CHNOSZ for thermodynamic calculations in aqueous geochemistry.
6.12: Problems
This page contains a comprehensive set of chemistry problems related to equilibrium constants, redox reactions, solubility, acid-base equilibrium, buffer solutions, and complexation reactions. It starts with deriving equilibrium constant expressions for given chemical reactions, analyzing the favorability of reactions using ladder diagrams, and calculating potentials for redox systems.
6.13: Additional Resources
The page provides a comprehensive list of references addressing various aspects of equilibrium chemistry. Topics covered include experimental determination of equilibrium constants, the impact of ionic strength, solubility products, and buffer capacity. Historical perspectives on the field are also offered. Additionally, the list encompasses instructional strategies, simulations for teaching, and critiques of conventional approaches to equilibrium concepts.
6.14: Chapter Summary and Key Terms
The chapter discusses analytical chemistry as the application of chemistry to analyze samples, focusing on using chemical reactivity to dissolve samples, separate analytes, transform analytes, or provide a signal. Key reactions include precipitation, acid-base, metal-ligand complexation, and oxidation-reduction. It also covers equilibrium concepts, such as Le Ch??telier's principle, and solutions like buffers, using equilibrium constants, ladder diagrams, and activity coefficients.

Support Center

How can we help?