15: Acid-Base Equilibria - A More Detailed Look
- Page ID
- 478930
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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 this chapter, we delve into the fascinating world of acids and bases, focusing on the dissociation of weak acids and bases and the use of ICE tables to solve equilibrium problems to quantify the concentrations of acids and bases in solutions. Understanding these fundamental concepts is crucial for mastering more complex topics in chemistry.
We highlight the work of Dr. Frank Wania. He is a renowned expert in environmental chemistry, particularly in the study of the environmental fate and transport of organic contaminants. His research is pivotal in understanding how chemicals interact with and move through the environment, which often involves investigating the dissociation properties of various compounds. Dr. Wania's work emphasizes the importance of basic physical-chemical properties, such as vapor pressure and the air/water Henry’s law constant. These properties are essential for predicting how contaminants will behave in different environmental contexts, mirroring the principles you will explore in this chapter about the dissociation of acids and bases. Through a combination of fieldwork, laboratory experimentation, and model simulations, Dr. Wania's research provides valuable insights into contaminant enrichment processes and helps inform strategies for mitigating environmental pollution.
Read more about Dr. Frank Wania and his research here: https://www.utsc.utoronto.ca/labs/wania/
As you study the principles of weak acids and bases and their dissociation, consider how these concepts are applied in real-world environmental chemistry scenarios. Dr. Wania’s investigations exemplify the critical role that a thorough understanding of chemical properties plays in addressing environmental challenges.
- 15.1: Introduction
- Acid-base chemistry revolves around proton (H⁺) transfer between donors (acids) and acceptors (bases) in water. These reversible reactions establish equilibria governing diverse phenomena, from sinkhole formation to oxygen transport in blood. This chapter delves deeper into the equilibrium principles underpinning this essential reaction class, highlighting its fundamental and applied contexts.
- 15.2: Arrhenius Theory- A Brief Review
- Acids and bases were historically identified by properties like sour taste, color changes in indicators (e.g., litmus), and reactions with metals/carbonates. Key milestones include Davy's recognition of hydrogen as acids essential component, and Arrhenius' definition, acids increase H⁺ ions (forming H₃O⁺) in water, while bases increase OH⁻ ions. This explained conductivity (acids/bases as electrolytes) and neutralization. However, the Arrhenius model is limited to aqueous solutions.
- 15.3: Brønsted-Lowry Theory of Acids and Bases
- The Brønsted-Lowry theory defines acids as proton (H⁺) donors and bases as proton acceptors, expanding beyond Arrhenius' aqueous limitations. Acid-base reactions involve proton transfer, forming conjugate pairs, acid/conjugate base and base/conjugate acid. Amphiprotic species act as both acids and bases. Water undergoes autoionization, governed by the ion-product constant. This enables equilibrium calculations for hydronium/hydroxide concentrations and explains acid-base behaviors.
- 15.4: Self-Ionization of Water and the pH Scale
- An equilibrium exists in pure water between its protonated and unprotonated form. In this section we examine this equilibrium in more detail.
- 15.5: Strong Acids and Strong Bases
- Acids and bases that are completely ionized when dissolved in water are called strong acids and strong bases There are only a few strong acids and bases, and everyone should know their names and properties. These acids are often used in industry and everyday life. The concentrations of acids and bases are often expressed in terms of pH, and as an educated person, you should have the skill to convert concentrations into pH and pOH. The pH is an indication of the hydrogen ion concentration, [H+].
- 15.6: Weak Acids and Weak Bases
- When acids or bases do not completely ionize, they are classified as weak acids or bases. In this chapter we will analyze the dissociate of these weak acids and bases using equilibrium constants.
- 15.7: Polyprotic Acids
- An acid that contains more than one ionizable proton is a polyprotic acid. The protons of these acids ionize in steps. The differences in the acid ionization constants for the successive ionizations of the protons in a polyprotic acid usually vary by roughly five orders of magnitude. As long as the difference between the successive values of Ka of the acid is greater than about a factor of 20, it is appropriate to break down the calculations of the concentrations sequentially.
- 15.8: Ions as Acids and Bases
- Salt solutions exhibit neutral, acidic, or basic properties based on their ions' acid-base behavior. Anions of weak acids hydrolyze water to produce basic solutions (pH > 7). Cations of weak bases yield acidic solutions (pH < 7) by releasing H⁺. Small, highly charged metal ions acidify water via Lewis acid hydration, weakening O–H bonds. Salts with equally matched conjugate pairs remain neutral. Hydrolysis is fundamentally an acid-base reaction governed by ion strength and charge density.
- 15.9: Molecular Structure and Acid-Base Behavior
- Inductive effects and charge delocalization significantly influence the acidity or basicity of a compound. The acid–base strength of a molecule depends strongly on its structure. The weaker the A–H or B–H+ bond, the more likely it is to dissociate to form an \(H^+\) ion. In addition, any factor that stabilizes the lone pair on the conjugate base favors the dissociation of \(H^+\), making the conjugate acid a stronger acid.
- 15.10: Lewis Acids and Bases
- Lewis Acid/Base Theory helps us understand acid/base behaviour in the most general terms by focusing on how electrons are involved in acid/base reactions.
- 15.11: Common-Ion Effect in Acid-Base Equilibria
- The common-ion effect suppresses dissociation or solubility when adding an ion already in equilibrium, per Le Châtelier's principle. In acid-base systems, this reduces ionization (↓ H⁺ in acetic acid with added acetate). In solubility systems, it decreases dissolved ions (↓ Pb²⁺ in PbCl₂ with added Cl⁻), often triggering precipitation. Equilibrium constants (K) remain unchanged, but the reaction quotient (Q) shifts position. This principle supports pH control in buffers and mineral formation.
- 15.12: Buffer Solutions
- Buffers are encountered not only in the chemistry lab but also in many biological systems as well. This section examines fundamental concepts of buffer solutions and shows how to analyze them quantitatively from an equilibrium perspective.
- 15.13: Acid-Base Indicators
- Acid-base indicators are weak organic acids or bases that exhibit distinct color changes based on pH due to shifts in their dissociation equilibrium: H I n ⇌ H + + I n − HIn HX + +InX − where HIn (protonated form) and In⁻ (deprotonated form) have different colors. The color-change interval spans ~2 pH units centered at the indicator's pKₐ (where [HIn] = [In⁻]).
- 15.14: Neutralization Reactions and Titration Curves
- Titration curves plot pH against titrant volume. For strong acid-strong base titrations, the curve is S-shaped with a sharp pH change at the equivalence point (pH = 7). Weak acid-strong base titrations show shallower initial pH changes, a higher equivalence point pH (>7), and a midpoint where pH = pKₐ. Weak bases exhibit inverse behavior. Polyprotic acids display multiple steps corresponding to each proton dissociation. Indicator selection requires matching pKᵢₙ to the equivalence point pH.
- 15.15: Solutions of Salts of Polyprotic Acids
- Salts from polyprotic acids hydrolyze stepwise, with pH determined by ion behavior. Strong acid/strong base salts are neutral. For acidic/basic ions, acidic cations lower pH. Basic anions raise pH. Amphoteric ions require comparing Ka (acid strength) vs. Kb (base strength = Kw/Ka of conjugate acid). If Kb > Ka, solution is basic; use ICE tables for quantification. General rule: polyprotic acid salts yield pH > 7.
- 15.18: Key Equations
- These are key equations. You should know how and when to use them. To practice their applications, check out 15.20.
- 15.19: Summary
- This section provides a brief summary of all the sections covered. It is important to understand that this should not be seen as a summary that will help you understand everything; it will only get you familiar with big concepts.
- 15.20: Exercises
- These end-of-chapter exercises are for practice to help you both understand concepts from the course and for practice for tests and exams.