Acid-Base Equilibria

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Acid-Base Equilibria

• Acids and bases are some of the more commonly
  encountered chemicals
• Acids and Bases control composition of blood and
  cell fluids, affect flavors, involved in
  digestion
• Bases used in house hold cleaners (NH3-based
  cleansers)
• Acid rain is an environmental problem
• Acids and bases are involved in reactions that
  produce polymers, synthetic fibers, dyes.

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Arrhenius Acid Base

• Acid produces H in aqueous solution
• Base produces OH- in aqueous solution
• HCl(aq) ? H (aq) Cl- (aq)
• NaOH(aq) ? Na (aq) OH- (aq)
• Acid base neutralization H(aq) OH-(aq) ?
  H2O(l)
• However, an H cannot exist by itself in water

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Brønsted-Lowry Acids and Bases

• Acid proton donor
• Base proton acceptor
• H PROTON since the H consists of 1 proton and
  0 electron
• HCN(aq) NH3(aq) ? NH4(aq) CN-(aq)
• acid base
• The H is transferred from HCN to NH3
• HCN is said to have an acidic H, a hydrogen that
  can be donated as a H

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• HCl has an acidic H, but by itself cannot act
  as an acid
• However HCl(aq)
• HCl (aq) H2O (l) ? H3O (aq) Cl- (aq)

H3O hydronium ion
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• HCN - hydrogen cyanide
• HCN(aq) H2O(l) ? H3O(aq) CN-(aq)
• Only a fraction of HCN donate their H to H2O
• HCN is a weak acid
• At equilibrium there is both CN- and
  un-dissociated HCN

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• In the Brønsted-Lowry theory
• a strong acid is fully deprotonated in solution
• HCl (aq) H2O (l) ? H3O (aq) Cl- (aq)
• a weak acid is only partially deprotonated in
  solution
• HCN(aq) H2O(l) ? H3O(aq) CN-(aq)
• Typically the solvent is water, but not
  necessarily.
• An acid that is strong in water, may be weak in
  another solvent

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• Brønsted-Lowry Base
• A proton acceptor. In most cases the molecule
  possesses a lone pair of electrons to which a H
  can bond to.
• Example Oxide, O2-
• O2- (aq) H2O(l) ? 2 OH- (aq)
• Strong base since all O2- (aq)
• forms OH- (aq)

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• NH3 a Brønsted base. The lone pairs on N in NH3
  can bond with a H.
• NH3 (aq) H2O (l) ? NH4 (aq) OH- (aq)
• NH3(aq) is a weak base at equilibrium both
  undissociated NH3 (aq) and NH4 (aq) exist.
• A strong base is completely protonated in
  solution
• O2- (aq) H2O(l) ? 2 OH- (aq)
• A weak base is partially protonated in solution
• NH3 (aq) H2O (l) ? NH4 (aq) OH- (aq)
• Strength depends on solvent

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Solvent Leveling Since all strong acids are
completely de-protonated in water (behave as
though they were solutions of H3O) strong acids
are leveled in water To compare acidity of
acids that are strong acids in water, need to use
a solvent in which the acidity of the acids
differ Strong bases are leveled in water in the
same way as strong bases.
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• Arrhenius definition restricted to water as a
  solvent
• However Brønsted-Lowry theory includes
  non-aqueous solvents
• CH3COOH (l) NH3 (l) ? CH3COO- (am) NH4(am)
• am - denotes a species dissolved in ammonia
• Brønsted-Lowry includes acid/base in the absence
  of solvent
• Protons can be transferred in the gas phase
• HCl(g) NH3(g) ? NH4Cl(s)
• Acid-base reaction does not have to involve the
  solvent
• HCN(aq) NH3(aq) ? NH4(aq) CN-(aq)

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Conjugate Acids Bases

• HCN(aq) H2O(l) ? H3O(aq) CN-(aq)
• acid conjugate base
• CN- (aq) is the conjugate base of HCN
• Brønsted-Lowry acids form conjugate bases
• Acid -----------gt conjugate base

donates H
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• Brønsted-Lowry bases form conjugate acids
• NH3 (aq) is the base NH4 (aq) is the conjugate
  acid
• NH3 (aq) H2O (l) ? NH4 (aq) OH (aq)
• base -------------gt conjugate acid

accepts H
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An acid is a proton donor and a base is a proton
acceptor. The conjugate base of an acid is the
base formed when the acid has donated a
proton. The conjugate acid of a base is the acid
that forms when the base has accepted a proton.
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Lewis Acids Bases

• A Lewis base donates a lone pair of electrons
• A Lewis acid accepts a lone pair of electrons
• Lewis acids/bases are a broader definition than
  the Brønsted-Lowry definition
• H is an electron pair acceptor a Lewis acid

Soluble metal oxides are strong bases
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• NH3 H2O ? NH4 OH-
• base acid

Reactions between electron deficient and
electron-rich molecules BF3(g) NH3(g) ?
F3B - NH3 (s) Lewis acid Lewis base
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• B-N bond is called a coordinate covalent bond
  formed by the coordination of an electron-pair
  donor to an electron pair acceptor

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Amphoterism

• H2O acts as both an acid and a base - amphoteric
• H2O(l) H2O(l) ? H3O (aq) OH- (aq)
• OH- conjugate base of H2O
• H3O conjugate acid of H2O
• HCO3- is amphoteric
• HCO3- (aq) H2O(l) ? H3O (aq) CO32- (aq)
• HCO3- (aq) H2O(l) ? H2CO3 (aq) OH- (aq)

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• Water is amphiprotic - both an acid and a base
• When one molecule transfers a proton to another
  molecule of the same kind - autoprotolysis or
  autoionization
• 2 H2O (l) ? H3O(aq) OH- (aq)
• An O-H bond is strong the fraction of protons
  transferred is very small.

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• Calculate the equilibrium constant for the
  autoionization of H2O(l)
• 2 H2O (l) ? H3O(aq) OH- (aq)
• Kw H3O(aq) OH- (aq)
• DGro DGfo(H3O(aq)) DGfo(OH-(aq)) - 2
  DGfo(H2O(l))
• 79.89 kJ/mol
• DGro - R T ln Kw
• Kw 1.0 x 10-14 at 298 K

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• Kw 1.0 x 10-14 at 298 K
• Kw H3O(aq) OH- (aq)
• H3O(aq) OH- (aq) 1.0 x 10-14
• Kw is an equilibrium constant the product of the
  concentrations of H3O and OH- is always equal to
  Kw.
• In pure water H3O(aq) OH- (aq) 1.0 x
  10-7 M at 298 K

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If the concentration of OH-(aq) in increased,
then H3O(aq) decreases to maintain Kw.
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What are the molarities of H3O and OH- in 0.0030
M Ba(OH)2 at 25oC? Ba(OH)2 (aq) ? Ba2 (aq)
2 OH- (aq) Molarity of OH- (aq) 0.0060
M H3O (aq) Kw/OH- (aq) 1.7 x 10-12 M
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pH Scale

• The concentration of H3O can vary over many
  orders of magnitude
• A log scale allows a compact description of the
  H3O concentration.

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• pH - log H3O
• H3O 10- pH mol/L
• For pure water at 25oC
• pH - log (1.0 x 10-7) 7.00
• For a change in pH by 1, H3O concentration
  changes by 10
• Higher pH, lower H3O concentration
• pH of pure water is 7
• pH of an acidic solution is less than 7
• pH of a basic solution is greater than 7

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