# 6.14: Chapter Summary and Key Terms


## Chapter Summary

Analytical chemistry is more than a collection of techniques; it is the application of chemistry to the analysis of samples. As we will see in later chapters, almost all analytical methods use chemical reactivity to accomplish one or more of the following: dissolve a sample, separate analytes from interferents, transform an analyte into a more useful form, or provide a signal. Equilibrium chemistry and thermodynamics provide us with a means for predicting which reactions are likely to be favorable.

The most important types of reactions are precipitation reactions, acid–base reactions, metal‐ligand complexation reactions, and oxidation–reduction reactions. In a precipitation reaction two or more soluble species combine to produce an insoluble precipitate, which we characterize using a solubility product.

An acid–base reaction occurs when an acid donates a proton to a base. The reaction’s equilibrium position is described using either an acid dissociation constant, Ka, or a base dissociation constant, Kb. The product of Ka and Kb for an acid and its conjugate base is the dissociation constant for water, Kw.

When a ligand donates one or more pairs of electron to a metal ion, the result is a metal–ligand complex. Two types of equilibrium constants are used to describe metal–ligand complexation: stepwise formation constants and overall formation constants. There are two stepwise formation constants for the metal–ligand complex ML2, each of which describes the addition of one ligand; thus, K1 represents the addition of the first ligand to M, and K2 represents the addition of the second ligand to ML. Alternatively, we can use a cumulative, or overall formation constant, $$\beta_2$$, for the metal–ligand complex ML2, in which both ligands are added to M.

In an oxidation–reduction reaction, one of the reactants is oxidized and another reactant is reduced. Instead of using an equilibrium constants to characterize an oxidation–reduction reactions, we use the potential, positive values of which indicate a favorable reaction. The Nernst equation relates this potential to the concentrations of reactants and products.

Le Châtelier’s principle provides a means for predicting how a system at equilibrium responds to a change in conditions. If we apply a stress to a system at equilibrium—by adding a reactant or product, by adding a reagent that reacts with a reactant or product, or by changing the volume—the system will respond by moving in the direction that relieves the stress.

You should be able to describe a system at equilibrium both qualitatively and quantitatively. You can develop a rigorous solution to an equilibrium problem by combining equilibrium constant expressions with appropriate mass balance and charge balance equations. Using this systematic approach, you can solve some quite complicated equilibrium problems. If a less rigorous answer is acceptable, then a ladder diagram may help you estimate the equilibrium system’s composition.

Solutions that contain relatively similar amounts of a weak acid and its conjugate base experience only a small change in pH upon the addition of a small amount of strong acid or of strong base. We call these solutions buffers. A buffer can also be formed using a metal and its metal–ligand complex, or an oxidizing agent and its conjugate reducing agent. Both the systematic approach to solving equilibrium problems and ladder diagrams are useful tools for characterizing buffers.

A quantitative solution to an equilibrium problem may give an answer that does not agree with experimental results if we do not consider the effect of ionic strength. The true, thermodynamic equilibrium constant is a function of activities, a, not concentrations. A species’ activity is related to its molar concentration by an activity coefficient, $$\gamma$$. Activity coefficients are estimated using the extended Debye‐Hückel equation, making possible a more rigorous treatment of equilibria.

## Key Terms

 acid activity coefficient base dissociation constant charge balance equation dissociation constant equilibrium formation constant Henderson–Hasselbalch equation Le Châtelier’s principle metal–ligand complex Nernst equation pH scale precipitate reduction steady state acid dissociation constant amphiprotic buffer common ion effect enthalpy equilibrium constant Gibb’s free energy ionic strength ligand method of successive approximations oxidation polyprotic redox reaction standard‐state stepwise formation constant activity base buffer capacity cumulative formation constant entropy extended Debye‐Hückel equation half‐reaction ladder diagram mass balance equation monoprotic oxidizing agent potential reducing agent standard potential solubility product

This page titled 6.14: Chapter Summary and Key Terms is shared under a CC BY-NC-SA license and was authored, remixed, and/or curated by David Harvey.