Acid Base Equilibria

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Acid Base Equilibria
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Theres acid in my rain

• Normal rain has a pH of approximately 5.6
• Due to industrial processes and automobiles, rain
  with a pH below 5.6 may result. This is
  considered acid rain
• Acid rain has different effects in different
  geographic areas and upon different organisms

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Acid Rain effects

• Snails die at pH lt 6.0
• Perch and pike die at pH lt 5.0
• Eel and brook trout die at pH lt 4.5
• Poisonous metals such as Al, Hg, and Pb are more
  soluble in acidic water

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Natural Buffers

• Much of the soil and rock in Alberta contains
  carbonates, which can neutralize acids
• 2 H3O(aq) CaCO3(s)
• Ca2(aq) CO2(g) 3 H2O(l)
• Much of the Canadian Shield throughout Manitoba,
  Ontario and Quebec is granite, which does not
  contain carbonates. Acid rain in these areas is
  more detrimental to wildlife as a result.

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Efficiency - is it a bad thing?

• The reaction of nitrogen with oxygen does not
  readily occur at room temperature
• N2(g) O2(g) ?H 2 NO(g)
• Kc 4.7x10-31 at 298 K
• At 1100 K (the operating temperature of an
  engine), Kc 6x10-9 at 298 K
• So, we should lower the operating temperature of
  engines to reduce the production of NO(g), right?

• WRONG! At lower temperatures, fuel burns less
• effeciently so it is not practical to do so.

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Empirical Properties of acids and bases

• Acids
• Taste sour.
• Turns blue litmus red.
• pH less than 7.
• Neutralizes bases.
• Reacts with active metals to produce hydrogen
  gas.
• Feels itchy

• Bases
• Tastes bitter.
• Turns red litmus blue.
• pH greater than 7.
• Neutralizes acids.
• Feels slippery.

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Arrhenius definitions of acids and bases.

• Acids dissociate to produce hydrogen ions in
  solution.

• Bases dissociate to produce hydroxide ions in
  solution.

H(aq)
OH-(aq)
eg. HCl(g) H(aq) Cl-(aq)
eg. NaOH(s) Na(aq) OH-(aq)
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Criticisms of the Arrhenius Definitions

• Hydrogen-containing polyatomic ions - predicted
  to be neutral
• eg. NaHCO3 Na HCO3 -
• is, in fact, basic
• eg. NaHSO4 Na HSO4 -
• is, in fact, acidic

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Criticisms of the Arrhenius Definitions

• 2. Non-metallic oxides - predicted to be
  neutral
• SO2 (g) SO2 (aq)
• are, in fact, acidic.

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Criticisms of the Arrhenius Definitions

• 3. Metallic oxides - predicted to be neutral
• eg. MgO(s) Mg2(aq) O2-(aq)
• are, in fact, basic

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Criticisms of the Arrhenius Definitions

• 4. Transition metal salts are predicted to be
  neutral
• eg. Cu(NO3)2(s) Cu2(aq) NO3-(aq)
• yet many are, in fact, acidic

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Criticisms of the Arrhenius Definitions

• The Killer
• 5. There is no experimental evidence for the
  existence of the hydrogen ion in solution.

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Revised Arrhenius definitions of acids and bases.

• Acids dissociate or ionize to produce hydronium
  ions in solution.

• Bases dissociate or ionize to produce hydroxide
  ions in solution.

OH-(aq)
H(aq)
HCl(g) H2O(l) H3O(aq) Cl-(aq)
NH3(g) H2O(l) NH4(aq) OH-(aq)
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• 1.
• HCO3-(aq) H2O(l) H2CO3 (aq) OH-(aq)
• HSO4- (aq) H2O(l) SO42- (aq) H2O(l)
• 2.
• SO2(g) H2O(l) H2SO3(aq)
• H2SO3(aq) H2O(l) H3O(aq) HSO3-(aq)

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• 3.
• MgO (s) Mg2 (aq) O2- (aq)
• O2-(aq) H2O(l) 2 OH- (aq)

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• 4.
• Cu(NO3)(s) 6 H2O(l)
• Cu(H2O)62(aq) 2 NO3-(aq)
• Cu(H2O)62 (aq) H2O(l)
• Cu(H2O)5(OH)(aq) H3O(aq)

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• 5.
• How can you define something, based upon
  something that does not exist?

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Brønsted - Lowry Definitions

• Acid
• Proton donor

• Base
• Proton acceptor

HCl(g) H2O(l) H3O(aq) Cl-(aq)
NH3(g) H2O(l) NH4(aq) OH-(aq)
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Consider the reaction below


H

O

H
O

CH3
C


O
H


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What is pH?

• A measurement of the concentration of H3O ions
  (H ions)

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pH, pOH, H3O, and OH-

• Kw HOH- 1 x 10-14
• pH -logH
• pOH -logOH-

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pH Scale
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Concentrated/Dilute vs. Strong/Weak

• Recall that the term concentrated refers to the
  relative amount of a solute present in a given
  volume of solution. A more concentrated solution
  has more moles per litre than does a diluted
  solution.
• In terms of acids and bases, the strength refers
  to the degree of dissociation in solution. An
  acid which dissociates 100 is termed strong
  an acid which dissociates to a lesser extent is
  deemed weak.

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pH of Strong Acids

• Since a strong acid dissociates 100, its pH may
  be determined easily from its concentration
• eg. What is the pH of a 0.012 mol/L HCl(aq)
  solution?
• HCl(aq) H2O(l) Cl-(aq) H3O(aq)
• pH -log H3O(aq) -log HCl(aq) 1.92

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pH of Weak Acids

• Since Weak acids do not dissociate 100, the
  concentration of hydronium ions is NOT the
  concentration of the original acid.
• You cannot simply take the log of the original
  acid concentration to get the pH.
• You must calculate the extent of dissociation as
  shown in the following example

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• eg. What is the pH of a 0.012 mol/L solution
  of ethanoic acid?
• Since ethanoic acid is a weak acid, we need to
  use an ICE table
• CH3COOH(aq) H2O (l) CH3COO- (aq) H3O (aq)
• I 0.012 0 0
• C -x x x
• E 0.012-x x x

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• Now we can use the Ka value from the data booklet
  to complete the calculation
• Ka 1.8x10-5 H3OCH3COO-
• CH3COOH
• (x)(x) (x)(x)
• (0.012-x) (0.012)
• x 4.6x10-4 H3O CH3COOH

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• Now, we can calculate the pH
• pH -logH3O 3.33
• Compare this to the value of 1.92 obtained for
  HCl (a strong acid) at the same concentration.
• This is a characteristic of weak acids They will
  always have a higher pH than a strong acid at the
  same concentration. The weaker the acid, the
  higher the pH.

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Percent Dissociation

• The percent dissociation for an acid is simply a
  comparison of the dissociated form to the
  original concentration.
• In the preceding example, the dissociated form is
  CH3COO-.
• The percent dissociation is
• 4.6x10-4
• 0.012

x 100 3.8
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Predicting Brønsted-Lowry Acid-Base (BLAB)
Reactions

• The possible acid-base reaction that may occur in
  a reaction mixture may be predicted in a very
  similar manner to that used in redox reaction
  predicting, as the next few slides show.

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Five Step Method of Predicting Acid Base Reactions

• List all species. (Review on next slide).
• Identify all the possible acids (A - left side of
  acid/base table) and bases (B - right side of
  acid-base table) .
• Identify the strongest acid (SA - highest on the
  left) and the strongest base (SB - lowest on the
  right).
• Predict the products by transferring a proton
  (H) from the SA to the SB.
• Predict the position of the equilibrium.

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Listing Species in Acid Base reactions.

• Do not dissociate
• All insoluble ionic compounds. .(BaSO4(aq))
• Weak acids. (CH3COOH(aq))
• Molecular compounds. (C12H22O11 (aq) )
• Polyatomic species. (MnO4- (aq) )

• Dissociate
• All ionic soluble compounds.(NaNO3(aq))
• All strong acids into hydronium and anion.
  (HCl(aq) is written as H3O (aq) and Cl- (aq) )

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• eg. What is the expected reaction when solutions
  of sodium nitrate and potassium hydrogen
  carbonate are mixed?
• Na NO3- K HCO3- H2O

A
B
B
A
B
SB
SA
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• So, the predicted reaction is
• HCO3-(aq) HCO3-(aq) H2CO3(aq) CO32-(aq)
• Since the SA is below the SB, the reactants are
  favoured at equilibrium

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Amphoterism

• In the previous example, both water and the
  hydrogen carbonate ion were labelled as possible
  acids or bases. Such a substance is said to be
  amphoteric or amphiprotic. It has the ability to
  act as an acid in one situation and as a base in
  another.

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Buffers

• A buffer is a solution that is resistant to
  changes in pH when small amounts of acid or base
  are added.
• A buffer usually consists of similar
  concentrations of a weak acid and its conjugate
  base.
• There is a limit to the amount of stress a buffer
  can handle before drastic changes in pH occur.
  This is referred to as the buffering capacity.
• Generally, buffers with higher concentrations of
  the weak acid and its base have greater buffering
  capacity.

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Biological Buffers

• The pH of blood and other bodily fluids must be
  controlled with certain very narrow ranges to
  avoid negatively impacting cells
• The two major buffers in the body are
• H2CO3 / HCO3- and H2PO4- / HPO42-

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How Buffers Work

• We have already learned that the acid base
  reaction that occurs is between the strongest
  acid and the strongest base.
• Since buffers have both an acid and a base
  present, they can handle either an acid or base
  being added.

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How Buffers Work

• If an acid is added to the H2CO3 / HCO3-
  equilibrium, the following reaction happens
• HCO3- H3O H2CO3 H2O
• If a base is added
• H2CO3 OH- HCO3- H2O
• In both cases, the stress is relieved so the pH
  remains near the original level.

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pH of Buffers

• The pH of a buffer is easily calculated using an
  ICE table, in a manner very similar to that used
  for weak acids.
• The only major difference is that buffers already
  have an initial conjugate base concentration, so
  the shift from original conditions to equilibrium
  is even less than for weak acids.

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pH of Buffers

• For example, consider the H2CO3 / HCO3- buffer.
  If the solution is prepared to be 1.0 mol/L for
  each of the two components, what is the pH?

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pH of Buffers

• H2CO3 H2O HCO3- H3O
• I 1.0 1.0 0
• C -x x x
• E 1.0-x 1.0-x 1.0x 1.0x

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• Now, using the Ka value for HCO3-, we can
  calculate the equilibrium concentrations and thus
  the pH
• Ka 4.5x10-7 (1.0x)(x)
• (1.0-x)
• By neglecting x in both the numerator and
  denominator, we simplify the math to give
• Ka 4.5x10-7 (1.0)(x)
• (1.0)

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• Which gives
• x 4.5x10-7 H3O
• And pH 6.35
• If the pH range needs to be adjusted downward, a
  new buffer in which there is a slightly greater
  proportion of the base, can be prepared.