CHEM 1112 Acid Base Equilibria

Part I

Aqueous species

• species “dissolve” in water

• in some cases molecules remain intact, and interact with water through non-covalent interactions

• in some cases, the molecules no longer associate with each other and are surrounded by water molecules “hydration shell”

Ionic reactions

• ionic species dissociate completley in aqueous solution, and are therefore more realistically represented as dissociated ions

  • consider the reaction between KOH and HCl in water.

• this is not the proper form, rather…

• (above is before the spectator ions are eliminated)

Spectator Ions

• ions whose presence is required to maintain charge neutrality

• neither chemically nor physically changed by the progress of the reaction

  • identical form on both sides of the arrow

Net Ionic Equation

• Complete ionic equation written out, with the spectator ions eliminated

• easier to see which species are changed during the reaction

Solubility Rules

• All common salts of group 1 elements are soluble

• All ammonium ions (NH₄+) are soluble

• All common nitrates (NO₃-) are soluble

• Binary compounds of Group 17 (NOT F) with metals are soluble

• Strong bases are the hydroxides of Group 1 and 2 metals (NOT Be)

Strong base

• completley dissolves into ions when dissolved in water, releasing OH- ions

• Hydroxides of Group 1 and 2 metals (NOT Be)

  • LiOH

  • NaOH

  • KOH
    Ca(OH)₂

  • Sr(OH)₂

  • Ba(OH)₂

Strong acid

• completely ionizes into ions when dissolved in water, releasing H+ ions

  • HClO₄

  • HCl

  • HBr

  • HI

  • HNO₃

  • H₂SO₄

Arrhenius definition (Acid)

• a substance that produces H+ ions in aqueous solutions

• a proton (H+ ion) donor

Arrhenius definition (base)

• a substance that produces OH- ions in aqueous solutions

• a proton (H+) acceptor

Lewis definition (acid)

• electron pair acceptor

Lewis definition (base)

• electron pair donor

Conjugate base of an acid

• the base formed when an acid donates a single proton

• H₂O + HA ⇋ (H₃O+) + A-

  • HA is the acid

  • A- is the conjugate base

Conjugate acid of a base

• the acid formed by the acceptance of a single proton by a base

• (H₃O+) + B- ⇋ HB + H₂O

  • B- is the base

  • HB is the conjugate acid

Amphoteric

• can act as either an acid or a base

• water holds this property

  • 

Autoionize

• molecule can spontaneously form ions without external influences

Ionization of water

• water dissociates to form H+ and OH- ions

• dissociation constant (K_w) for this 10^-14 at 25^oC

Monoprotic acid

• some Bronsted acids are monoprotic, meaning they only have one acidic proton

• consider acetic acid (CH₃COOH), only one proton on the carboxyl group can be donated

  • 

Diprotic acid

• some Bronsted acids are diprotic, meaning they have two acid protons.

• consider sulfuric acid (H₂SO₄)…

pH scale

• used to express the hydronium ion concentration

• the acidity of a solution is expressed as a pH value where pH= -log[H+]

  • neutral solution: [H+] = [OH-]

  • basic solution: [H+] < [OH-]

  • acidic solution: [H+] > [OH-]

• more acidic the solution, the lower its pH

• a change in pH of one unit reflects a tenfold change in the [H+]

Monoprotic strong acids

• dissociate completely to donate their protons to water, so that there will be virtually no undissociated acid left in solution

  • HClO₄

  • HCl

  • HBr

  • HI

  • HNO₃

  • H₂SO₄

Strong bases

• completely dissociate in solution to form hydroxide ions

  • LiOH

  • NaOH

  • KOH

  • Ca(OH)₂

  • Sr(OH)₂

  • Ba(OH)₂

Strong vs weak acids

• strong acids dissociate completely to donate their protons to water

  • the [H+] = initial acid concentration

• weak acids do NOT dissociate completely to donate all of their protons to water.

  • the [H₃O+] =/= weak acid concentration

Acid ionization constant (K_a)

• K_a represents the equilibrium constant for the dissociate of an acid

  • indicates the strength of an acid

• the stronger the acid, the larger the K_a value

• K_a = [H₃O+][A-] / [HA]

% ionization of a weak acid

• % ionization = [H₃O+]_eqb’m / [HA]_initial

Weak bases

• do NOT dissociate completely in water

  • the concentration of HO- in a weak base solution is not equal to the concentration of the weak base

  • K_b = [BH+][HO-] / [B]

• the stronger the base, the larger the value of K_b

Sources of H+ in solutions

• Ionization of the acid/base

  • 

• Autoionization of water

  • usually insignificant

  • 

Oxyacids

• acids that contains an inner atom bonded to a variable # of oxygen atoms and acidic OH groups

• general formula: Y-OH

Amines

• large group of weak bases

• many other weak bases are derivatives of ammonia called amines, where some of the bonds to H have been replaced with bonds to other atoms

• lone pair can accept H+

• amines that have their free lone pair can act as a base (accept a proton)

Acid strength

• recall: the stronger the acid, the larger the K_a

  • K_a is related to ∆G^o

  • ∆G^o = -RTlnK_a

  • ∆G = ∆H - T∆S

• qualitatively, would not expect big difference in ∆S_r for different HA

  • all have same change in # of species, all species are aq

• differences must be due to different ∆H

  • the H+ is the same for all acid dissociates

  • qualitatively, the differences must be due to the relative differences between HA and A-

• thus, qualitatively, we can understand the magnitudes of ∆G and thus K_a by considering the relative stabilities of the acid and its conjugate base

Factors affecting relative stabilities of HA and A-

• electrostatic factors

• ability to conjugate base to accomodate the negative charge

  • resonance

  • size

  • other EN groups

effect of charge (acid strength)

• affects the ability to donate protons (opposite charges attract)

  • H₃O is better able to donate proton to a base than H₂O is; H₃O is a stronger acid than H₂O

  • HO- does not function as an acid in aq solutions

Acidity trends (across)

• as you move across a period, and thus increase the electronegativity of a species, acidity increases

  • acid strength increases with increasing # of EN groups attached

• increased polarity of species-H bond

• increased stability of the product

Acidity trend (down)

• increased size of the species as you move down the periodic table helps to better accomodate negative charge, and thus increases acidity

Carboxylic acids

• oxyacids

• organic (carbon) based

• weak acids