General Chemistry II 2302102 Acid and Base Equilibria [email protected]...
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Transcript of General Chemistry II 2302102 Acid and Base Equilibria [email protected]...
General Chemistry II2302102
Acid and Base Equilibria
Lecture 1
Acids and Bases - 3 Lectures
• Autoionization of Water and pH
• Defining Acids and Bases
• Interaction of Acids and Bases with Water
• Buffer Solutions
• Acid-Base Titrations
Outline - 5 Subtopics
By the end of this lecture AND completion of the set problems, you should be able to:
• Understand strong and weak electrolytes, Kw and the autoionization equilibrium in water, definition of the pH scale and its relationship to [H3O
+] and [OH-].
• Know the Arrhenius and Brønsted definitions of acids and bases.
• Understand and know examples of conjugate acid-base pairs.
• Be familiar with the ionization reactions for strong and weak acids, know examples of typical monoprotic, diprotic and triprotic acids.
• Understand the common ion effect in acid ionization.
Objectives - Lecture 1
Acids and Bases
Acids & Bases
Gold mining frequently uses the base, cyanide (CN-), in the extraction process. Tailings dams sometimes have high concentrations of CN-. The conjugate acid of CN-, HCN, is extremely toxic.
“Acid mine drainage” (AMD) is naturally produced by the exposure of sulfide ores to water. This process is greatly exacerbated by mining the sulfide ores (for Cu, Pb, Zn etc). AMD results in streams with elevated H+ concentrations and pH values in the range 1-3.
A gold mining tailings dam leaks into an AMD-affected stream - would you evacuate?
Would you evacuate?
Acids & Bases - 2 Lectures
Introduction to acids and basesStrong & Weak AcidsConjugate acid-base pairsCommon Ion Effect
BasesBuffersIndicatorsTitrations
• Strong Acids• Weak Acids
Outline
Acids & Bases - Lecture 1
By the end of this lecture AND completion of the set problems, you should be able to:
• Define BrØnsted acids and bases• Calculate [H+], [OH-] and pH• Distinguish between strong & weak acids• Calculate equilibrium concentrations of acids &
bases using acidity constants• Identify conjugate acid-bases pairs• Determine the effect of adding a common ion on
equilibrium concentrations
Objectives
Acids & Bases - Why are we interested?
Acids - solvents (dissolve other materials)- ores, food (stomach contains HCl)
Bases - solvents - cleaners like “Draino”, bleach
Large Range of Industrial Processes:
Acids - wine, beer, citrus fruits, vinegar, coffeeFoods:
Blood - pH 7.3-7.5 falls below 6.8 will be fatal (acidosis)Most bodily functions under ‘circumneutral’ conditions
Physiology:
‘Natural’ Erosion of Limestone Caves“acid rain” - dissolved H2SO4 & HNO3 in upper atmosphere
- rainfall runoff into acid lakes - devoid of life
Environment:
Acids & Bases
1. ACID Species which can donate a hydrogen ion (H+).
2. BASE Species which can accept a hydrogen ion.
3. AMPHOTERIC SPECIES Species which can act as both an acid and a
base. e.g. HCO3- (CO3
2-, H2CO3), HSO4-
BrØnsted-Lowry Definitions:
• Hydrogen Ion
What is the Hydrogen Ion?• Hydrogen Atom Atomic weight = 1
1 proton + 1 electron
• Hydrogen Ion in Water
Represented as:
H+(aq) or H3O+
Atomic weight = 1
1 proton
Acid-Base ReactionIn fact a proton-transfer reaction in which the proton is transferred from the acid to the base.
Acidity - pH
pH convenient measure of H3O+ concentration
H O
concentration of H O M
pH
pH
3
3
10
10
e.g. pH = 8.5
H O
concentration of H O M
38
9
39
10
3 16 10
3 16 10
.5
.
.
Acidity measure of the H3O+ concentration
AQUEOUS SYSTEMS
2H2O() H3O+(aq) + OH-(aq)
Equilibrium between H3O+ and OH-:
at equilibrium [H3O+][OH-] = Kw
at 25 °C Kw = 1.00 x 10-14
(c.f. Ksp = [Ag+][Cl-])
AQUEOUS SYSTEMS
2H2O() H3O+(aq) + OH-(aq)
Kw = [H3O+][OH-]
at 25 °C Kw = 1.00 x 10-14
Pure water at 25 °C. If 2z mol/L of water react:
[H3O+] = z and [OH-] = z
z2 = 1.00 x 10-14 z = 1.00 x 10-7
and pH = 7.00
DEFINITIONS
“acid” pH < 7.00 [H3O+(aq)] > [OH-(aq)]
“basic” pH > 7.00 [H3O+(aq)] < [OH-(aq)]
“neutral” pH = 7.00 [H3O+(aq)] = [OH-(aq)]
Acid-Base Reaction: Proton-Transfer between an Acid and a Base (Water)
• According to the Brönsted-Lowry concept, when HCl gas is dissolved in water to form the solution of hydrochloric acid, a proton-transfer reaction occurs:
H2O (l) + HCl (g) H3O+ (aq) + Cl- (aq)
BaseBase(Proton(Proton
Acceptor)Acceptor)HydroniumHydronium
IonIonA hydrated proton = HA hydrated proton = H++ (aq) = H (aq) = H33OO++
AcidAcid(Proton(ProtonDonor)Donor)
ClClOOHH ::
HH
++ HH :: HH OO HH
HH:: ClCl ::
--++
++::
:::: :: ::
::
ACIDS
Note HA may be a molecule anion
or cation
HCl aq HCOOH aq
HCO aq
NH aq
( ), ( )
( )
( )
3
4
Reaction between acids and water
H2O() + HA(aq) H3O+(aq) + A-(aq)
Equilibrium constant Ka
At equilibrium
A aq H O aq
HA aqKa
( ) ( )
( )
3
ACIDSReaction between acids and water
Equilibrium constant Ka
H2O() + HA(aq) H3O+(aq) + A-(aq)
pK K
so that K
a a
apK
a
log10
10
Usually tabulated as pKa:
Note - this is exactly the same relationship as between [H+] and pH (pH = - log10 [H+])
Strength of Acids
• Strong versus Weak acids
• The strength of an acid is related to the position of the equilibrium above
• It is NOT related to the corrosive ability (this often causes confusion)
• As we shall see, HF is a weak acid, yet it is one of the most corrosive acids known.
H2O() + HA(aq) H3O+(aq) + A-(aq)
STRONG ACID
Position of equilibrium almost completely to the right (acid is almost totally ionized)
e.g. 0.1 M HCl(aq) [H3O+] = 0.1 [Cl-] = 0.1 [HCl] not known ( < 10-10)
pKa < -10
Reaction between acids and water
H2O() + HA(aq) H3O+(aq) + A-(aq)
Acids are ElectrolytesNon
No ionsin solution
Strong
Many ionsin solution -
Weak
Few ionsin solution -
Strong Acids Weak Acids
WEAK ACIDS
1. Definition of weak acid Position of equilibrium lies to the left. Only
a small fraction of the acid reacts with the water and is ionized. i.e. Acid is weakly ionized
Ka small < 10-3
pKa large > 3
Reaction between acids and water
H2O() + HA(aq) H3O+(aq) + A-(aq)
WEAK ACIDS
initial 0.5 M 0 0equilibrium (0.5 - x) M x M x M
2. Dissociation in water
H2O() + HA(aq) H3O+(aq) + A-(aq)
eg. 0.5 M HF(aq) Ka = 6.8 x 10-4
If x mol/L react then we have
H2O() + HF(aq) H3O+(aq) + F-(aq)
initial 0.5 M 0 0equilibrium (0.5 - x) M x M x M
At equilibrium:
MHF
M
M
OH
F
48.0][
108.1][
108.1][
23
2
H2O() + HF(aq) H3O+(aq) + F-(aq)
So HF is ca. 4% ionized
WEAK ACIDS
Diprotic & Triprotic Acids• Acids that we’ve considered thus far have
been monoprotic (donate 1 proton) e.g. HCl, HNO3
• Other acids can donate 2 or 3 protons - diprotic or triprotic acids e.g. H2SO4, H3PO4H2O() + H2A(aq) H3O+(aq) + HA-(aq)
]A[
][ ][
HHAOH
2
31
K a
H2O() + HA-(aq) H3O+(aq) + A2-(aq)
][
][ ][
HA
AOH 23
2
K a
Diprotic & Triprotic Acids
]A[
][ ][
HHAOH
2
31
K a
][
][ ][
HA
AOH 23
2
K a
Acid Formula Ka1 pKa1 Ka2 pKa2 Ka3 pKa3
Sulfuric H2SO4 ca. 100 ca. -2 1.2 x 10-21.92
Oxalic H2C2O4 5.9 x 10-21.23 6.4 x 10-5
4.19
Phosphoric H3PO4 7.52 x 10-32.12 6.23 x 10-8
7.21 2.2 x 10-1312.67
Successive pKas increase in magnitude
- removal of successive protons results in weaker acids
Conjugate Acid-Base Pairs
- Reverse reactionH2O (l) + NH3 (aq) NH4
+ (aq) + OH- (aq)
Consider the proton-transfer reaction:– Forward reaction
H2O (l) + NH3 (aq) NH4+ (aq) + OH- (aq)
Acid(ProtonDonor)
Base(Proton
Acceptor)
Acid(ProtonDonor)
BaseBase(Proton(Proton
Acceptor)Acceptor)
Conjugate Acid-Base Pairs
H2O (l) + NH3 (aq) NH4+ (aq) + OH- (aq)
Acid 1 Base 1 Acid 2 Base 2
Conjugate Acid-Base Pair
Conjugate Acid-Base Pair
For the general equation
B + HA HB+ + A-
Conjugate Acid-Base Pairs
Acid(Proton (Proton Donor)Donor)
Base(Proton (Proton
Acceptor)Acceptor)
Acid(Proton (Proton Donor)Donor)
Base(Proton (Proton
Acceptor)Acceptor)
The acid HB+ results from the base B gaining a proton. HB+-B is a conjugate acid-base pair.
The base A- results from the acid HA losing a proton. HA-A- is a conjugate acid-base pair, too.
Conjugate Acid-Base Pairs
• Any two substances that differ by one proton
are a conjugate acid-base pair
• Can write the formula of the conjugate base of
any acid simply by removing the proton:If HCO3
- is the acid, CO32- is the conjugate base
If HCO3- is the base, H2CO3 is the conjugate acid
(HCO3- is amphoteric)
Conjugate Acid-Base Pairs
• The conjugate (base) of a strong acid is a weak base
• The conjugate (base) of a weak acid is a strong base
• The conjugate (acid) of a strong base is a weak acid
• The conjugate (acid) of a weak base is a strong acid
Direction of Acid-Base Reactions
NH4+ is a stronger acid than H2O, and
OH- is a stronger base than NH3
So reaction proceeds spontaneously to the left
• Acid-Base Reactions proceed spontaneously with the strongest acid and strongest base forming the weakest acid and the weakest base
Returning to:H2O (l) + NH3 (aq) NH4
+ (aq) + OH- (aq)
The Common Ion Effect
Recall that with sparingly soluble salts:The presence of an ion in solution which is common to the electrolyte will decrease the solubility
(adding KF to a solution of PbF2 reduced the Pb2+ concentration)
The presence of an ion in solution which is common to the weak acid will suppress it’s ionization (decrease the concentration of the free ion)
The Common Ion Effect
initial 0.5 M 0 0equilibrium (0.5 - x) M x M x M
H2O() + HF(aq) H3O+(aq) + F-(aq)
At equilibrium, [F-] = [H3O+] = 1.8 x 10-2 Mso pH = 1.74
What happens if we add 0.05 M NaF?
initial 0.5 M 0 0.05equilibrium (0.5 - x) M x M 0.05 + x M
H2O() + HF(aq) H3O+(aq) + F-(aq)
The Common Ion Effect
initial 0.5 M 0 0.05equilibrium (0.5 - x) M x M 0.05 + x M
H2O() + HF(aq) H3O+(aq) + F-(aq)
So (0.05 + x) x = (0.5 - x) 6.8 x 10-4
x = 6.00 x 10-3 M So [H3O+] = 6.00 x 10-3 M
Thus pH = 2.22 c.f. 1.74 without the NaF
Acids and Bases - End of Lecture 1
After studying this lecture and the set problems, you should be able to:
• Understand strong and weak electrolytes, Kw and the autoionization equilibrium in water, definition of the pH scale and its relationship to [H3O
+] and [OH-].
• Know the Arrhenius and Brønsted definitions of acids and bases.
• Understand and know examples of conjugate acid-base pairs.
• Be familiar with the ionization reactions for strong and weak acids, know examples of typical monoprotic, diprotic and triprotic acids.
• Understand the common ion effect in acid ionization.
Objectives Covered in Lecture 1