Acids and Bases

Acids, Bases and Equilibrium When an acid is dissolved in water, the H + ion (proton) produced by the acid combines with water to produce the hydronium ion, H 3 O + When an acid is dissolved in water, the H + ion (proton) produced by the acid combines with water to produce the hydronium ion, H 3 O + HCl and other strong electrolytes ionize completely in water HCl and other strong electrolytes ionize completely in water Weak acids like acetic acid ionize only to a very small extent in water Weak acids like acetic acid ionize only to a very small extent in water

Acids, Bases and Equilibrium The equilibrium concept is used to describe to what extent an acid or base ionizes in water The equilibrium concept is used to describe to what extent an acid or base ionizes in water Ionization constants: K>1 indicates a strong acid or base, K 1 indicates a strong acid or base, K<1 refers to a weak acid or base

Strong Electrolytes Strong Acids – HCl, HBr, HI, HNO 3, HClO 4, H 2 SO 4 (for first H + only) Strong Acids – HCl, HBr, HI, HNO 3, HClO 4, H 2 SO 4 (for first H + only) Strong Bases – LiOH, NaOH, KOH, RbOH, CsOH, Sr(OH) 2, Ba(OH) 2 Strong Bases – LiOH, NaOH, KOH, RbOH, CsOH, Sr(OH) 2, Ba(OH) 2 All others are probably weak All others are probably weak

Arrhenius Acids and Bases An acid is a substance that, when dissolved in water, increases the concentration of hydrogen ions, H+, in the water An acid is a substance that, when dissolved in water, increases the concentration of hydrogen ions, H+, in the water A base is a substance that, when dissolved in water increases the concentration of hydroxide ion, OH-, in the water A base is a substance that, when dissolved in water increases the concentration of hydroxide ion, OH-, in the water

*Bronsted-Lowry Acids and Bases An acid is any substance that can donate a proton to any other substance An acid is any substance that can donate a proton to any other substance Examples of Bronsted acids – molecular compounds (HNO 3 ), cations (NH 4 + ), hydrated metals (Fe(H 2 O) 6 3+ ), or anions (H 2 PO 4 - ) Examples of Bronsted acids – molecular compounds (HNO 3 ), cations (NH 4 + ), hydrated metals (Fe(H 2 O) 6 3+ ), or anions (H 2 PO 4 - ) Theory is not restricted to compounds in water Theory is not restricted to compounds in water

Bronsted-Lowry Acids and Bases A Bronsted base is a substance that can accept a proton from any other substance A Bronsted base is a substance that can accept a proton from any other substance Bronsted bases can be molecular compound (NH 3 ), anions (CO 3 2- ), or cations (Al(H 2 O) 5 (OH) 2+ ) Bronsted bases can be molecular compound (NH 3 ), anions (CO 3 2- ), or cations (Al(H 2 O) 5 (OH) 2+ )

Polyprotic Acids + Bases Acid Form Amphiprotic Form Base Form H2SH2SH2SH2S HS - S 2- H 3 PO 4 H 2 PO 4 - HPO 4 2- PO 4 3- H 2 CO 3 HCO 3 - CO 3 2- H2C2O4H2C2O4H2C2O4H2C2O4 HC 2 O 4 - C 2 O 4 2-

Conjugate Acid-Base Pairs A pair of compounds or ions that differ by the presence of one H+ ion is called a conjugate acid-base pair A pair of compounds or ions that differ by the presence of one H+ ion is called a conjugate acid-base pair Every reaction between a Bronsted acid and Bronsted base involves H+ transfer and has two conjugate acid- base pairs Every reaction between a Bronsted acid and Bronsted base involves H+ transfer and has two conjugate acid- base pairs

*Conjugate Acid-Base Pairs Acid 1 Base 2 Base 1 Acid 2 HCl + H 2 O  Cl - + H 3 O + HCO H 2 O  CO H 3 O + CH 3 CO 2 H + H 2 O  CH 3 CO H 3 O + H2OH2OH2OH2O + NH 3  OH - + NH 4 + H2OH2OH2OH2O + CO 3 2-  OH - + HCO 3 - H2OH2OH2OH2O + H 2 O  OH - + H 3 O +

*Water and the pH scale Water autoionizes to a small extent, producing low concentrations of H 3 O + and OH - ions (water conducts electricity) Water autoionizes to a small extent, producing low concentrations of H 3 O + and OH - ions (water conducts electricity) The equilibrium for autoionization of water lies far to the left The equilibrium for autoionization of water lies far to the left At 25 o C, Kw=1.0 x (the ionization constant for water) At 25 o C, Kw=1.0 x (the ionization constant for water)

*Water and the pH scale Kw increases with temperature because the autoionization of water is endothermic Kw increases with temperature because the autoionization of water is endothermic Kw is valid in pure water and any aqueous solution Kw is valid in pure water and any aqueous solution In pure water and dilute aqueous solutions, the concentration of water is considered to be constant at 55.5 M In pure water and dilute aqueous solutions, the concentration of water is considered to be constant at 55.5 M

*Water and the pH scale [H 3 O + ] = [OH - ] = 1.0 x M in pure water, a neutral solution [H 3 O + ] = [OH - ] = 1.0 x M in pure water, a neutral solution In acidic solution, [H 3 O + ]>[OH - ] In acidic solution, [H 3 O + ]>[OH - ] In basic solution, [H 3 O + ]<[OH - ] In basic solution, [H 3 O + ]<[OH - ]

*Water and the pH scale pH = -log[H 3 O + ] pH = -log[H 3 O + ] pOH = -log[OH - ] pOH = -log[OH - ] pKw = pH + pOH = pKw = pH + pOH = 14.00

Relationship between hydronium and hydroxide ion concentrations, pH and pOH

*Equilibrium Constants for Acid and Bases The strength of acids and bases of the same concentration can be compared by measuring pH The strength of acids and bases of the same concentration can be compared by measuring pH The relative strength of an acid can be expressed with an equilibrium constant The relative strength of an acid can be expressed with an equilibrium constant K a = [H 3 O + ][A - ]/[HA] K a = [H 3 O + ][A - ]/[HA]

Equilibrium Constants for Acid and Bases K b = [BH + ][OH - ]/[B] K b = [BH + ][OH - ]/[B] Weak acid, K a 2, small [H 3 O + ] Weak acid, K a 2, small [H 3 O + ] Weak base, Kb<1, pH<12, small [OH - ] Weak base, Kb<1, pH<12, small [OH - ] A large value of K indicates ionization products are strongly favored; small K value indicates reactants are favored A large value of K indicates ionization products are strongly favored; small K value indicates reactants are favored

Equilibrium Constants for Acid and Bases Weak acids have strong conjugate bases; small Ka corresponds with large Kb Weak acids have strong conjugate bases; small Ka corresponds with large Kb Consider the acids and conjugate bases in table 15.2 on page 668 Consider the acids and conjugate bases in table 15.2 on page 668 Notice trends: as acid strength declines in a series, the relative conjugate base strength increases Notice trends: as acid strength declines in a series, the relative conjugate base strength increases

*Equilibrium Constants for Acid and Bases Use Tables 15.3 on page 671, 15.4 on page 679, and 15.5 on page 683 Use Tables 15.3 on page 671, 15.4 on page 679, and 15.5 on page 683 Which is the stronger acid, H 2 SO 4 or H 2 SO 3 ? Which is the stronger acid, H 2 SO 4 or H 2 SO 3 ? Is benzoic acid stronger or weaker than acetic acid? Is benzoic acid stronger or weaker than acetic acid? Which has the stronger conjugate base, acetic acid or formic acid? Which has the stronger conjugate base, acetic acid or formic acid? Which is the stronger base, ammonia or methylamine? Which is the stronger base, ammonia or methylamine? Which has the stronger conjugate acid, ammonia or methylamine? Which has the stronger conjugate acid, ammonia or methylamine?

Calculations with Equilibrium Constants The principles of the equilibria can be applied to aqueous solutions of weak acids and bases. The equilbrium constants Ka and Kb can be determined if the concentrations of the various species present in the solution are known. These are often determined by measuring pH. The principles of the equilibria can be applied to aqueous solutions of weak acids and bases. The equilbrium constants Ka and Kb can be determined if the concentrations of the various species present in the solution are known. These are often determined by measuring pH.

*Calculations with Equilibrium Constants If the acid or base is weak, and the initial concentration of acid (or base) is at least 100x Ka (or Kb), then the approximation that If the acid or base is weak, and the initial concentration of acid (or base) is at least 100x Ka (or Kb), then the approximation that [acid] initial = [acid] equilibrium is valid. Otherwise the quadratic equation must be solved. [acid] initial = [acid] equilibrium is valid. Otherwise the quadratic equation must be solved.

Calculations with Equilibrium Constants (ex) For the weak acid: (ex) For the weak acid: HA (aq) + H 2 O (l)  H 3 O + (aq) + A- (aq) Ka = [H 3 O + ][A-]/[HA] and because [H 3 O + ] = [A-], Ka = [H 3 O + ] 2 /[HA] The assumption that [HA] equil = [HA] initial is valid if [HA] initial > 100 x Ka

*Calculations with Equilibrium Constants Calculating a Ka value from a measured pH Calculating a Ka value from a measured pH A solution prepared from mol butanoic acid dissolved in sufficient water to give 1.0 L of solution has a pH of Determine Ka for butanoic acid. The acid ionizes according to the balanced equation: A solution prepared from mol butanoic acid dissolved in sufficient water to give 1.0 L of solution has a pH of Determine Ka for butanoic acid. The acid ionizes according to the balanced equation: CH 3 CH 2 CH 2 CO 2 H + H 2 O  H 3 O + + H 3 CH 2 CH 2 CO 2 - CH 3 CH 2 CH 2 CO 2 H + H 2 O  H 3 O + + H 3 CH 2 CH 2 CO 2 -

CH 3 CH 2 CH 2 CO 2 H + H 2 O  H 3 O + + H 3 CH 2 CH 2 CO 2 - Calculate initial molarity of butanoic acid Calculate initial molarity of butanoic acid Calculate equilibrium molarity of hydronium ion from pH Calculate equilibrium molarity of hydronium ion from pH Construct ice table Construct ice table Write equilibrium constant expression and substitute in values from ice table Write equilibrium constant expression and substitute in values from ice table Calculate Ka Calculate Ka

*Calculating Equilibrium Concentrations and pH from Ka What are the equilibrium concentrations of acetic acid, acetate ion, and hydronium ion for a 0.10 M solution of acetic acid (Ka = 1.8 x )? What is the pH of the solution? What are the equilibrium concentrations of acetic acid, acetate ion, and hydronium ion for a 0.10 M solution of acetic acid (Ka = 1.8 x )? What is the pH of the solution? Write chemical equation for ionization of acetic acid in water Write chemical equation for ionization of acetic acid in water ICE it up ICE it up Write equilibrium constant expression, substitute values from ice table, solve for x Write equilibrium constant expression, substitute values from ice table, solve for x Use x to find equilibrium concentrations Use x to find equilibrium concentrations Use concentration of hydronium ion to solve for pH Use concentration of hydronium ion to solve for pH

*Calculating the pH of a salt solution What is the pH of a M solution of sodium acetate? What is the pH of a M solution of sodium acetate? Write equation for ionization reaction of acetate ion in water (it is a base!) Write equation for ionization reaction of acetate ion in water (it is a base!) Put it on ICE Put it on ICE Write equilibrium constant expression and substitute ICE values Write equilibrium constant expression and substitute ICE values Solve for x, solve for concentration of hydroxide ion Solve for x, solve for concentration of hydroxide ion Solve for concentration of hydronium ion, solve for pH Solve for concentration of hydronium ion, solve for pH

*Calculating the pH after the reaction of an acid with a base Calculate the pH after mixing 15 mL of 0.12 M acetic acid with 15 mL of 0.12 M NaOH. What are the major species in solution at equilibrium (besides water) and what are their concentrations? Calculate the pH after mixing 15 mL of 0.12 M acetic acid with 15 mL of 0.12 M NaOH. What are the major species in solution at equilibrium (besides water) and what are their concentrations? Write balanced equations Write balanced equations Stoichiometry problem to solve for “ initial ” concentration of acetate anion Stoichiometry problem to solve for “ initial ” concentration of acetate anion ICE, equilibrium constant express Kb, pOH and such ICE, equilibrium constant express Kb, pOH and such

*Aqueous Solutions of Salts CationAnion pH of Solution From strong base (Na+) From strong acid (Cl-) =7 neutral From strong base (K+) From weak acid (CH 3 CO 2 -) >7 basic From weak base (NH 4 +) From strong acid (Cl-) <7 acidic From any weak base (BH+) From any weak acid (A-) Depends on relative strengths of acid and base

*For each of the following salts in water, predict whether the pH will be greater than, less than, or equal to 7 KBr KBr NH 4 NO 3 NH 4 NO 3 AlCl 3 AlCl 3 Na 2 HPO 4 Na 2 HPO 4

*Polyprotic acids and bases The pH of many inorganic polyprotic acids depends primarily on the hydronium ion generated in the first ionization step The pH of many inorganic polyprotic acids depends primarily on the hydronium ion generated in the first ionization step Each successive loss of a proton is about more difficult than the previous step Each successive loss of a proton is about more difficult than the previous step The hydronium ion produced in the second step can be neglected; calculate using the k of the first ionization The hydronium ion produced in the second step can be neglected; calculate using the k of the first ionization

*Molecular Structure, Bonding, and Acid Strength In a series of acids, as bond strength decreases, acid strength increases In a series of acids, as bond strength decreases, acid strength increases Adjacent electronegative atoms that pull electrons from a hydrogen increase the strength of an acid (inductive effect) Adjacent electronegative atoms that pull electrons from a hydrogen increase the strength of an acid (inductive effect) Acids that have resonance structures are stronger because they are more stable after they lose a proton than acids without resonance Acids that have resonance structures are stronger because they are more stable after they lose a proton than acids without resonance tml tml tml tml

*Lewis Acids and Bases A Lewis acid is a substance that can accept a pair of electrons from another atom to form a new bond A Lewis acid is a substance that can accept a pair of electrons from another atom to form a new bond A Lewis base is a substance that can donate a pair of electrons to another atom to form a new bond A Lewis base is a substance that can donate a pair of electrons to another atom to form a new bond This is also know as coordinate covalent chemistry This is also know as coordinate covalent chemistry

Lewis Acids and Bases A + B:  B:A A + B:  B:A acid + base  acid-base adduct acid + base  acid-base adduct The acid must have an empty orbital available or be able to make one available The acid must have an empty orbital available or be able to make one available A base must have a pair of nonbonding electrons A base must have a pair of nonbonding electrons The formation of hydronium ion is an example of this type of reaction The formation of hydronium ion is an example of this type of reaction

Lewis Acids and Bases Good Lewis bases include hydroxide ion and ammonia and water Good Lewis bases include hydroxide ion and ammonia and water Metal cations are good Lewis acids; transition metals form complexes known as coordination complexes Metal cations are good Lewis acids; transition metals form complexes known as coordination complexes Nonmetal oxides such as carbon dioxide are Lewis acids Nonmetal oxides such as carbon dioxide are Lewis acids Some metal hydroxides are amphoteric – acting like an acid in the presence of a base and a base in the presence of an acid (ex. Al(OH) 3 ) Some metal hydroxides are amphoteric – acting like an acid in the presence of a base and a base in the presence of an acid (ex. Al(OH) 3 )

* pKa = -log Ka * pKa = -log Ka A logarithmic scale is used to report and compare acid strengths A logarithmic scale is used to report and compare acid strengths The pKa value becomes smaller as the acid strength increases The pKa value becomes smaller as the acid strength increases KaKb = Kw for an acid and its conjugate base KaKb = Kw for an acid and its conjugate base Also, pKw = pKa + pKb Also, pKw = pKa + pKb

*Equilibrium Constants and Acid-Base Reactions All proton transfer reactions proceed from the stronger acid and base to the weaker acid and base. All proton transfer reactions proceed from the stronger acid and base to the weaker acid and base. Write the net ionic equation for the possible reaction between acetic acid, Ka = 1.8 x and sodium hydrogen sulfite, NaHSO 3, Kb = 8.3 x Does the equilibrium lie to the left or right? Write the net ionic equation for the possible reaction between acetic acid, Ka = 1.8 x and sodium hydrogen sulfite, NaHSO 3, Kb = 8.3 x Does the equilibrium lie to the left or right?

Types of Acid-Base Reactions Reaction of a Strong acid with a Strong Base Reaction of a Strong acid with a Strong Base Net ionic equation: Net ionic equation: H 3 O + + OH -  2H 2 O H 3 O + + OH -  2H 2 O K = 1/Kw = 1.0 x K = 1/Kw = 1.0 x Mixing equal molar quantities of a strong base with a strong acid produces a neutral solution (pH=7, at 25 o ) Mixing equal molar quantities of a strong base with a strong acid produces a neutral solution (pH=7, at 25 o )

Types of Acid-Base Reactions Unless both the acid and base involved in the neutralization reaction are strong, the pH of the solution that results will not be neutral. Unless both the acid and base involved in the neutralization reaction are strong, the pH of the solution that results will not be neutral.

Types of Acid-Base Reactions Reaction of a Weak Acid with a Strong Base Reaction of a Weak Acid with a Strong Base (ex.) net ionic equation: (ex.) net ionic equation: CH 3 CO 2 H + OH -  H 2 O + CH 3 CO 2 - CH 3 CO 2 H + OH -  H 2 O + CH 3 CO 2 - Adding these two reactions give the net ionic equation above; Adding these two reactions give the net ionic equation above; CH 3 CO 2 H + H 2 O  H 3 O + + CH 3 CO 2 - Ka=1.8x10 -5 CH 3 CO 2 H + H 2 O  H 3 O + + CH 3 CO 2 - Ka=1.8x10 -5 H 3 O + + OH-  2 H 2 O Ka=1/Kw=1.0x10 14 H 3 O + + OH-  2 H 2 O Ka=1/Kw=1.0x10 14 The K for the net reaction is the product of the two equilibrium constants K neut =Ka x 1/Kw=1.8x10 9 The K for the net reaction is the product of the two equilibrium constants K neut =Ka x 1/Kw=1.8x10 9

Types of Acid-Base Reactions Mixing equal molar quantities of a strong base with a weak acid produces a salt whose anion is the conjugate base of the weak acid. The solution is basic, with the pH depending on Kb for the anion. Mixing equal molar quantities of a strong base with a weak acid produces a salt whose anion is the conjugate base of the weak acid. The solution is basic, with the pH depending on Kb for the anion. In the example above, the salt that results from the reaction is acetate, the conjugate base of a weak acid. Therefore the solution will be basic at the equivalence point. In the example above, the salt that results from the reaction is acetate, the conjugate base of a weak acid. Therefore the solution will be basic at the equivalence point.

Types of Acid-Base Reactions Reaction of a Strong Acid with a Weak Base Reaction of a Strong Acid with a Weak Base (ex.) net ionic equation: (ex.) net ionic equation: H 3 O + + NH 3  NH H 2 O H 3 O + + NH 3  NH H 2 O Adding these two reactions give the net ionic equation above: Adding these two reactions give the net ionic equation above: NH 3 + H 2 O  NH 4 + OH- Kb = 1.8x10 -5 NH 3 + H 2 O  NH 4 + OH- Kb = 1.8x10 -5 H 3 O + + OH-  2 H 2 O K = 1/Kw = 1.0x10 14 H 3 O + + OH-  2 H 2 O K = 1/Kw = 1.0x10 14 The K for the net reaction is the product of the two equilibrium constants K = Kb x 1/Kw = 1.8x10 9 The K for the net reaction is the product of the two equilibrium constants K = Kb x 1/Kw = 1.8x10 9

Types of Acid-Base Reactions Mixing equal molar quantities of a strong acid and a weak base produces a salt whose cation is the conjugate acid of the weak base. The solution is acidic, with the pH depending on the Ka for the cation. Mixing equal molar quantities of a strong acid and a weak base produces a salt whose cation is the conjugate acid of the weak base. The solution is acidic, with the pH depending on the Ka for the cation. In the above example, the solution at the equivalence point contains the ammonium ion, the conjugate acid of a weak base, and the solution is acidic. In the above example, the solution at the equivalence point contains the ammonium ion, the conjugate acid of a weak base, and the solution is acidic.

Types of Acid-Base Reactions Reaction of a Weak Acid with a Weak Base Reaction of a Weak Acid with a Weak Base (ex.) net ionic equation: (ex.) net ionic equation: CH 3 CO 2 H + NH 3  NH CH 3 CO 2 - CH 3 CO 2 H + NH 3  NH CH 3 CO 2 - The reaction is product-favored because acetic acid is stronger than ammonium ion and ammonia is a stronger base than acetate ion (use table) The reaction is product-favored because acetic acid is stronger than ammonium ion and ammonia is a stronger base than acetate ion (use table) K = Ka x Kb / Kw (what is the K for this ex?) K = Ka x Kb / Kw (what is the K for this ex?)

Types of Acid-Base Reactions When a weak acid reacts with a weak base, the pH of the solution at the equivalence point depends upon which is the stronger, the acid or the base When a weak acid reacts with a weak base, the pH of the solution at the equivalence point depends upon which is the stronger, the acid or the base If equal molar solutions are mixed the resulting solution contains ammonium acetate; is it acidic or basic? (use table) If equal molar solutions are mixed the resulting solution contains ammonium acetate; is it acidic or basic? (use table)

Types of Acid-Base Reactions Mixing equal molar quantities of a weak acid and a weak base produces a salt whose cation is the conjugate acid of the weak base and whose anion is the conjugate base of the weak acid. The solution pH depends on the relative Ka and Kb values. Mixing equal molar quantities of a weak acid and a weak base produces a salt whose cation is the conjugate acid of the weak base and whose anion is the conjugate base of the weak acid. The solution pH depends on the relative Ka and Kb values.