13: Electrolytes, Acids and Bases
- Page ID
- 568792
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)- 13.1: Electrolytes
- Substances that dissolve in water to yield ions are called electrolytes. Electrolytes may be covalent compounds that chemically react with water to produce ions (for example, acids and bases), or they may be ionic compounds that dissociate to yield their constituent cations and anions, when dissolved. Dissolution of an ionic compound is facilitated by ion-dipole attractions between the ions of the compound and the polar water molecules.
- 13.2: Acids- Properties and Examples
- This page provides an overview of various acids, emphasizing their properties, common examples, and applications. It covers notable acids like hydrochloric, sulfuric, and nitric acids, as well as citric acid and acetylsalicylic acid (aspirin). The importance of acids in foods, beverages, and industrial uses is highlighted, showcasing their roles in flavoring, preservation, and medicinal applications.
- 13.3: Bases- Properties and Examples
- This page covers the properties and applications of bases, contrasting them with acids. It describes bases as electrolytes, which can be strong or weak, with a bitter taste and slippery texture. Common indicators reveal their presence through color changes. Notable bases like sodium hydroxide and magnesium hydroxide are highlighted for their use in household products and antacids, underscoring their importance in neutralizing acids and their widespread use in cleaning and health products.
- 13.4: Strong and Weak Acids and Bases
- This page defines strong and weak acids and bases, detailing their dissociation characteristics in solution. Strong acids and bases fully ionize, while weak ones only partially do so. It includes a table of common strong acids and bases, offers examples, and provides exercises for identifying their strength. Additionally, it highlights that salts from weak acids or bases can affect the acidity or basicity of a solution.
- 13.5: Arrhenius vs. Bronsted-Lowry Acid-Base Theory
- This page covers the definitions and classifications of acids and bases according to the Arrhenius and Brønsted-Lowry theories, distinguishing between proton donors and acceptors. It highlights the limitations of the Arrhenius theory and introduces the hydronium ion. Additionally, it elaborates on amphoteric species and conjugate acid-base pairs, illustrated with examples such as ammonia and acetic acid, emphasizing the role of proton transfer in acid-base equilibria.
- 13.6: Autodissociation of Water and Kw
- This page explains the autoionization of water, showing that water can act as both an acid and a base, forming hydronium (\(\ce{H3O^{+}}\)) and hydroxide ions (\(\ce{OH^{-}}\)). It details the equilibrium concentration of these ions and the relationship between them, emphasizing the constant \(K_w\) which varies with temperature.
- 13.7: The pH and pOH Scales
- This page covers the concepts of pH and pOH, explaining their definitions and calculations based on hydronium and hydroxide ion concentrations. It highlights that pH indicates acidity (pH < 7), neutrality (pH = 7), and basicity (pH > 7). The page includes the formula \(pK_w = pH + pOH\), walkthroughs for calculating pH and pOH, and stresses the importance of significant figures. Additionally, it features examples and practice exercises to reinforce understanding of these concepts.
- 13.8: Acid-Base Reactions
- This page explains acid-base neutralization reactions, detailing how acids and bases combine to form water and salts. It presents the net ionic equation by removing spectator ions and discusses how the pH of the solutions depends on the strengths of the reactants. It also covers acids reacting with metals to produce salts and hydrogen gas, along with bases reacting similarly. The page includes examples and exercises for further understanding.
- 13.9: Acid–Base Titrations
- This page covers the basics of acid-base titrations, detailing the procedure of adding a standard solution to an unknown until neutralization, indicated by a color change. It defines the equivalence point for complete reactions and explains how to calculate unknown concentrations. The importance of accurate measurement with a burette and suitable indicators for different titrations is discussed, supplemented with examples and exercises for practice.
- 13.10: Buffers- Solutions that Resist pH Change
- This page discusses buffers, which are solutions that maintain stable pH levels when acids or bases are introduced, composed of weak acids and their salts or weak bases with corresponding salts. Examples cited include acetic acid with sodium acetate and ammonia with ammonium chloride. It explains the reaction of buffers to strong acids and bases while noting their limited capacity for effectiveness. The importance of buffers in biological systems, especially in human blood, is also emphasized.