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 + HCl and other strong electrolytes ionize completely 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 Ionization constants: 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 Bases – LiOH, NaOH, KOH, RbOH, CsOH, Sr(OH) 2 , Ba(OH) 2 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 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 Examples of Bronsted acids – molecular compounds (HNO 3 ), cations (NH 4+ ), hydrated metals (Fe(H 2 O) 63+ ), or anions (H 2 PO 4 ) 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 Bronsted bases can be molecular compound (NH 3 ), anions (CO 32 ), or cations (Al(H 2 O) 5 (OH) 2+ )

Polyprotic Acids + Bases

Acid Form H 2 S H 3 PO 4 H 2 CO 3 H 2 C 2 O 4 Amphiprotic Form HS H 2 PO 4 HPO 42 HCO 3 HC 2 O 4 Base Form S 2 PO 43 CO 32 C 2 O 42-

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 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 HCl HCO 3 + H 2 O + H 2 O CH 3 CO 2 H + H 2 O H 2 O H 2 O H 2 O + NH + CO + H 2 3 3 O 2       Base 1 Cl CO 3 2 CH 3 CO 2 OH OH OH Acid 2 + H 3 O + + H 3 O + + H 3 O + + NH 4 + + HCO 3 + 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) The equilibrium for autoionization of water lies far to the left At 25 o C, Kw=1.0 x 10 -14 constant for water) (the ionization

Water and the pH scale

   Kw increases with temperature because the autoionization of water is endothermic 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

Water and the pH scale

   [H 3 O + ] = [OH ] = 1.0 x 10 -7 water, a neutral solution M in pure In acidic solution, [H 3 O + ]>[OH ] In basic solution, [H 3 O + ]<[OH ]

Water and the pH scale

   pH = -log[H 3 O + ] pOH = -log[OH ] 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 relative strength of an acid can be expressed with an equilibrium constant K a = [H 3 O + ][A ]/[HA]

Equilibrium Constants for Acid and Bases

    K b = [BH + ][OH ]/[B] Weak acid, K a <1, pH>2, small [H 3 O + ] 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

Equilibrium Constants for Acid and Bases

   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 Notice trends: as acid strength declines in a series, the relative conjugate base strength increases

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Equilibrium Constants for Acid and Bases

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 ?

Is benzoic acid stronger or weaker than acetic acid?

Which has the stronger conjugate base, acetic acid or formic acid?

Which is the stronger base, 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.

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 [acid] initial = [acid] equilibrium is valid. Otherwise the quadratic equation must be solved.

Calculations with Equilibrium Constants

 (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] valid if [HA] initial equil > 100 x Ka = [HA] initial is

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Calculations with Equilibrium Constants

Calculating a Ka value from a measured pH A solution prepared from 0.055 mol butanoic acid dissolved in sufficient water to give 1.0 L of solution has a pH of 2.72. 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

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H 3 O + + H 3 CH 2 CH 2 CO 2 -

     Calculate initial molarity of butanoic acid Calculate equilibrium molarity of hydronium ion from pH Construct ice table Write equilibrium constant expression and substitute in values from ice table 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 10 -5 )? What is the pH of the solution?

Write chemical equation for ionization of acetic acid in water ICE it up Write equilibrium constant expression, substitute values from ice table, solve for x Use x to find equilibrium concentrations Use concentration of hydronium ion to solve for pH

Calculating the pH of a salt solution

      What is the pH of a 0.015 M solution of sodium acetate?

Write equation for ionization reaction of acetate ion in water (it is a base!) Put it on ICE Write equilibrium constant expression and substitute ICE values Solve for x, solve for concentration of hydroxide ion Solve for concentration of hydronium ion, solve for pH

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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?

Write balanced equations   Stoichiometry problem to solve for “ initial ” concentration of acetate anion ICE, equilibrium constant express Kb, pOH and such

Polyprotic acids and bases

   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 10 4 -10 6 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

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Molecular Structure, Bonding, and Acid Strength

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) Acids that have resonance structures are stronger because they are more stable after they lose a proton than acids without resonance

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 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

Lewis Acids and Bases

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

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Lewis Acids and Bases

Good Lewis bases include hydroxide ion and ammonia and water Metal cations are good Lewis acids; transition metals form complexes known as coordination complexes 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 )

Acid and Base Properties of Some Ions in Aqueous Sol

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Neutral Basic Acidic Anions Cl Br I NO 3 ClO 4 Cations Li + Na + K + Ca 2+ Ba 2+ CH 3 CO 2 SO 4 2 HCO 2 CN OCl HPO 4 2 PO 4 3 NO 2 CO 3 2 S 2 HCO HS 3 SO 3 F 2 Al(H ions 2 O) 5 (OH) 2+ and analogous HSO 4 H 2 PO 4 HSO 3 Al(H NH 4 + 2 O) 6 3+ and hydrated transition metal cations

Acid and Base Properties of Some Ions in Aqueous Sol

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   Anions that are conjugate bases of strong acids are such weak bases that they have no effect on pH There are numerous basic anions that are conjugate bases of weak acids Acidic anions arise from polyprotic acids

Acid and Base Properties of Some Ions in Aqueous Sol

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  Alkali metal and alkaline earth metal cations have no effect on pH (conj. acid of strong base) All metal cations are hydrated in water, M(H 2 O) 6 n+ , but when M is +2 or +3 and particularly a transition metal, the ion acts as an acid

Aqueous Solutions of Salts

Cation From strong base (Na+) From strong base (K+) From weak base (NH 4 +) From any weak base (BH+) Anion From strong acid (Cl-) From weak acid (CH 3 CO 2 -) From strong acid (Cl-) From any weak acid (A-) pH of Solution =7 neutral >7 basic <7 acidic 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 NH 4 NO 3 AlCl 3 Na 2 HPO 4

     pKa = -log Ka A logarithmic scale is used to report and compare acid strengths The pKa value becomes smaller as the acid strength increases KaKb = Kw for an acid and its conjugate base 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.

Write the net ionic equation for the possible reaction between acetic acid and sodium hydrogen sulfate, NaHSO 4 . Does the equilibrium lie to the left or right?

Types of Acid-Base Reactions

     Reaction of a Strong acid with a Strong Base Net ionic equation: H 3 O + + OH  2H 2 O K = 1/Kw = 1.0 x 10 14 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.

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Types of Acid-Base Reactions

Reaction of a Weak Acid with a Strong Base (ex.) net ionic equation: CH 3 CO 2 H + OH  H 2 O + CH 3 CO 2 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 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

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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.

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 (ex.) net ionic equation: H 3 O + + NH 3  NH 4 + + H 2 O Adding these two reactions give the net ionic equation above: 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 The K for the net reaction is the product of the two equilibrium constants K = Kb x 1/Kw = 1.8x10

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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.

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 (ex.) net ionic equation: CH 3 CO 2 H + NH 3  NH 4 + + 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) 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 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.