16.S: Acid–Base Equilibria (Summary)
- Page ID
- 91292
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\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)16.1: Acids and Bases: A Brief Review
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- acids have sour taste and turn litmus paper red
- bases have a bitter taste and feel slippery
- Svante Arrhenius (1859-1927)
- Acids associated with H+ ions
- Bases associated with OH- ions
- Solution is acidic if there is more H+ than OH-
- Solution is basic if there is more OH- than H+
16.2: Brønsted–Lowry Acids and Bases
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- Arrhenius definition of acids and bases
- Acids when dissolved in water increase H+ concentration
- Bases when dissolved in water increase OH- concentration
- Arrhenius definition of acids and bases
16.4.1 Proton Transfer Reactions
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- Brq nsted-Lowry definition of acids an bases
- Acid is a proton donor
- Base is a proton acceptor
- Can be applied to non-aqueous solutions
- Brq nsted-Lowry acid must be able to lose a H+ ion
- Brq nsted-Lowry base must have a non bonding pair of electrons to bind to H+ ion
- Amphoteric - substance that can act as an acid or base
16.4.2 Conjugate Acid-Base Pairs
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- conjugate acid - product formed by adding a proton to base
- conjugate base - product formed by removal of a proton from acid
16.4.3 Related Strengths of Acids an Bases
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- the stronger the acid, the weaker the conjugate base
- the stronger the base, the weaker the conjugate acid
- equilibrium favors transfer of proton from stronger acid to stronger base
16.3: The Autoionization of Water
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- autoionization of water - dissociation of H2O molecules to H+ and OH- ions
- at room temperature only 1 out of 109 molecules are ionized
- exclude water from equilibrium expressions involving aqueous solutions
- ion-product constant
- kw = k[H2O] = [H+][OH-] = 1.0 x 10-14 (at 25° C)
- solution is neutral when [H+] = [OH-]
- solution is acidic when [H+] > [OH-]
- solution is basic when [H+] < [OH-]
16.3.1 The Proton in Water
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- H+ ion is a proton with no valence electrons
- H+ ion react with H2O molecule to form H3O+, hydronium ion
- H3O+ ion can bond with other H2O molecules to form hydrated hydrogen ions
- H+ and H3O+ used interchangeably
16.4: The pH Scale
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- concentration of [H+] expressed in terms of pH
- pH = -log [H+]
- acidic solutions
- [H+] > 1.0 x 10-7
- [OH-] < 1.0 x 10-7
- pH < 7.00
- neutral solutions
- [H+] = [OH-] = 1.0 x 10-7
- pH = 7
- basic solutions
- [H+] < 1.0 x 10-7
- [OH-] > 1.0 x 10-7
- pH > 7
Other "p" Series
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- pOH = -log [OH-]
- pH + pOH = -log Kw = 14.00
Measuring pH
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- pH meter
- has a pair of electrodes connected to a meter that measures in millivolts
- voltage generated when electrodes placed in solution, and is measured by meter
- red litmus paper for pH of 5 or lower
- blue litmus paper for pH of 8 or higher
- pH meter
16.5: Strong Acids and Bases
strong acids and bases are strong electrolytes
16.5.1 Strong Acids
strongest monoprotic acids
- HCl, HBr, HI, HNO3, HclO3, HclO4, and diprotic H2SO4
- For strong monoprotic acid concentration of [H+] equals the original concentration of the acid
16.5.2 Strong Bases
- most common strong bases are ionic hydroxides of alkali metals and the heavier alkaline-earth metals
- complete dissociation
16.6: Weak Acids
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- HA(aq) + H2O(l) ß à H3O+ + A-(aq)
- HA(aq) ß à H+(aq) + A-(aq)
- Ka = acid - dissociation constant
- The lager the Ka the stronger the acid
- Ka usually less than 10-3
16.6.1 Calculating pH for Solutions of Weak Acids
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- 1) write ionization equilibrium
- 2) write equilibrium expression
- 3) I.C.E. Table
- 4) substitute equilibrium concentrations into equilibrium expression
- in weak acids [H+] is small fraction of concentration of acid
- percent ionization depends on temperature, identity of acid and concentration
- as percent ionization decreases, concentration increases
16.6.2 Polyprotic Acids
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- more than one ionizable H atom
- easier to remove first proton than second
- acid dissociation constants are Ka1, Ka2, etc…
- Ka values usually differ by 103
16.7: Weak Bases
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- base-dissociation constant, Kb
- equilibrium at which base reacts with H2O to form a conjugate acid and OH-
- contain 1 or more lone pair of electrons
16.7.1 Types of Weak Bases
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- weak bases have NH3 and anions of weak acids
16.8: Relationship Between Ka and Kb
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- when two reactions are added together than equilibrium constant of third reaction is equal to the product of the equilibrium constants of the added reactions
- reaction 1 + reaction 2 = reaction 3
- K1 x K2 = K3
- Ka x Kb = [H+][OH-] = Kw
- Acid-dissociation constant times base-dissociation constant equals the ion-product constant for water
- Ka x Kb = Kw = 1.0 x 10-14
- pKa x pKb = pKw = 14; (pKa= -log Ka and pKb = -log Kb)
16.9: Acid-Base Properties of Salt Solutions
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- hydrolysis - ions reacting with water to produce H+ and OH- ions
- anions from weak acids react with water to produce OH- ions which is basic
- anions of strong acids are not basic and do not influence pH
- anions that have ionizable protons are amphoteric
- behavior depends on Ka and Kb
- all cations except those of alkali metals and heavier alkaline earth (Ca2+, Sr2+ and Ba2+) are weak acids in water
- alkali metal and alkaline earth cations do not hydrolyze
- do not affect pH
- strengths of acids and bases from salts
- 1) salts derived from strong acid and base
- no hydrolysis and solution has pH of 7
- 2) salts derived from strong base and weak acid
- strong conjugate base
- anion hydrolyzes and produces OH- ions
- cation does not hydrolyze
- pH greater than 7
- 3) salts derived from weak base and strong acids
- cation is strong conjugate acid
- cation hydrolyzes to produce H+
- anion does not hydrolyze
- solution has pH below 7
- 4) salts derived from weak acid and base
- both cation and anion hydrolyze
- pH depends on extent on hydrolysis of each ion]
16.10: Acid-Base Behavior and Chemical Structure
16.10.1 Factors that Affect Acid Strength
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- strength of acid depends on:
- 1) polarity of H-X bond
- 2) strength of H-X bond
- 3) stability of conjugate base, X-
- molecule will transfer proton if H-X bond is polarized
- in ionic hydrides H- acts as proton acceptor because of negative charge
- nonpolar bonds produce neither acidic or basic solutions
- strong bonds less easily dissociated that weak bonds
- the greater the stability of conjugate base, the stronger the acid]
- strength of acid depends on:
16.10.2 Binary Hydrides
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- metal hydrides are basic or have no acid-base properties in water
- nonmetal hydrides can be between having no acid-base properties to being acidic
- in each group of nonmetallic elements, acidity increases with increasing atomic number
- bond strengths decrease as central atom gets larger and overlap of orbitals get smaller
16.10.3 Oxyacids
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- Y-O-H bond
- Oxyacids - have OH bonded to central atom
- Base if bonded to a metal because pair of electrons shared between Y-O is completely transferred to O
- Ionic compound with OH- is formed
- When bonded to nonmetal the bond is covalent and compounds are acidic or neutral
- As electronegativity of Y increases , acidity also increases
- O-H bond becomes more polar
- Conjugate base usually an anion and stability increases as electronegativity of Y increases
- Relating acid strengths of oxyacids to electronegativity of Y and to number of groups attached to Y
- 1) same number of oxygen atoms, acid strength increases as electronegativity of central atom increases
- 2) same central atom Y, acid strength increases with increasing number of bonded oxygen atoms to central atom
- acidity increases as oxidation number of central atom increases
16.10.4 Carboxylic Acids
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- carboxyl group - COOH
- acidic behavior of carboxylic acids
- addition oxygen atom in carboxyl group draws density from O-H bond which increases the polarity
- conjugate base ion have resonance forms
- acidity increases as number of electronegative atoms in acid increases
16.11: Lewis Acids and Bases
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- Lewis acid - electron pair acceptor
- Lewis base - electron pair donor
- Any Brq nsted-Lowry is a Lewis base
- Lewis acids have molecules that have incomplete octets
16.11.1 Hydrolysis of Metal Ions
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- hydration - attraction of metal ions to water molecules
- metal ion acts as Lewis acid
- water molecule acts as Lewis base
- electron density drawn from oxygen atom to water molecule
- O-H bond becomes more polarized
- For hydrolysis reactions Ka increases with increasing charge and decreasing radius of ion
- hydration - attraction of metal ions to water molecules