Acid-Base Titrations
Introduction to Acids and Bases
Chapter 8
Name Acids Base
Definition Example Definition Example
Arrhenius An H-containing compound that increases H+ concentration in water
HClHNO3
HBrHFAcetic acid
An OH-containing compound that increases OH- concentration in water
KOHNaOH
Bronsted-Lowry
Proton donor Terminal alkynes
Proton acceptor NH3
Lewis Electron pair acceptor
Fe3+ and other transition metal cations
Electron pair donor
There are three (3) definitions of acids and bases: 1. Arrhenius, 2. Bronsted-Lowry and 3. Lewis.
There are three (3) definitions of acids and bases: 1. Arrhenius, 2. Bronsted-Lowry and 3. Lewis.
acidbase acid
acidbase conjugate acid
Conjugate acid-base pairs differ by only one proton.
base
conjugate base
There are three (3) definitions of acids and bases: 1. Arrhenius, 2. Bronsted-Lowry and 3. Lewis.
Lewis Acids:Any species that accepts electron pairs. (H+ CO2,
Mn+)
Arrhenius Bases:LiOH,
Mg(OH)2
Arrhenius Acids:
HCl, H2SO4
Bronsted Bases:H2O, NH3, HS-
Bronsted Acids:
H2O, HS-,
H2PO4-
Lewis Bases:Any species with
a lone pair
Acids and Bases undergo neutralization reactions.
NaOH + HCl NaCl + H2O
Na+ + OH– + H+ + Cl– Na+ + Cl– + H2O H+ + OH– H2O
Acids and Base strengths are defined by how much H3O+ or OH- is produced by a given concentration
Strong Acid Weak Acid
Acids and Base strengths are defined by how much H3O+ or OH- is produced by a given concentration
percent ionization =
Ionized acid concentration at equilibrium
Initial concentration of acidx 100%
For a monoprotic acid HA
[H+]eqx 100% [HA]0 = initial concentration
Percent ionization = [HA]0
Weak acids and bases are in equilibrium with their original species.
CH3COOH + H2O H3O+ +CH3COO–
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Kc =H3O
+[ ] CH3COO
−[ ]
H2O[ ] CH3COOH[ ]Kc << 1
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Kc H2O[ ] =H3O
+[ ] CH3COO
−[ ]
CH3COOH[ ]
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Ka =H3O
+[ ] CH3COO
−[ ]
CH3COOH[ ]
Weak acids and bases are in equilibrium with their original species.
NH3 + H2O NH4+
+OH–
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Kc =NH4
+[ ] OH
−[ ]
H2O[ ] NH3[ ]Kc << 1
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Kc H2O[ ] =NH4
+[ ] OH
−[ ]
NH3[ ]
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Kb =NH4
+[ ] OH
−[ ]
NH3[ ]
Ka and Kb tells us something about the relative acid/base strengths
Ka and Kb tells us something about the relative acid/base strengths
Water can act both as an acid and a base (AUTOIONIZATION)
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Kc =H3O
+[ ] OH
−[ ]
H2O[ ]2
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Kc H2O[ ]2
= H3O+
[ ] OH−
[ ]
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Kw = H3O+
[ ] OH−
[ ]
The ion-product constant (Kw) is the product of the concentration of H3O+ and OH– ions at a particular temperature
At 250CKw = [H3O+][OH-] = 1.0 x
10-14
The p-scale conveniently handles a wide range of concentrations
]log[ 3 OHpH
]log[ OHpOH
KwpKw log KapKa log
Solving weak acid ionization problems:
1. Identify the major species that can affect the pH.
• In most cases, you can ignore the autoionization of water. [H3O+] from water is negligible in comparison to [H3O+] from the weak acid.
• Ignore [OH-] because it is determined by [H+].
2. Use ICE to express the equilibrium concentrations in terms of single unknown x.
3. Write Ka in terms of equilibrium concentrations. Solve for x by the approximation method. If approximation is not valid, solve for x exactly.
4. Calculate concentrations of all species and/or pH of the solution.
ExercisesWhat is the pH of a 0.5 M HF solution (at 25°C,
Ka = 7.1 x 10-4)?
What is the pH of a 0.05 M HF solution?
What is the pH of a 0.122 M monoprotic acid whose Ka is 5.7 x 10-4?
Complete the following table:
Parameter
An NaOH solution
A 0.30 M HNO3
pH 9.52
pOH
[H3O+]
[OH-]
Complete the following table:
Parameter
An NaOH solution
A 0.30 M HNO3
pH 9.52 0.52
pOH 4.48 13.48
[H3O+] 3.0 x 10-
10M0.30 M
[OH-] 3.3 x 10-5
M3.3 x 10-14 M
Exercises
What is the pH, [H3O+], [OH-] of 7.52 x 10-4 M CsOH?
What is the pOH, [H3O+], [OH-] of 1.59 x10-3 M HClO4?
What is the [H3O+], [OH-] and pOH in a solution with a pH of 2.77
What is the [H3O+], [OH-] and pH in a solution with a pOH of 11.27
ExercisesChloroacetic acid has a pKa of 2.87. What are
[H3O+], pH, [ClCH2COOH], [ClCH2COO-] in 1.05 M [ClCH2COOH]
A 0.735 M of weak acid is 12.5% dissociated. Calculate [H3O+], pH, [OH-], pOH of solution. Calculate Ka of acid
BuffersChapter 9
Buffers solutions are solutions that resist changes in pH
Buffers contain appreciable amounts of a weak acid and its conjugate base
HA/ A–
NH4Cl/NH3
H3PO4/NaH2PO4
NH4SH/Na2SHCOOH/HCOOK
HBr/KBrH3IO3/Li2HIO3
NaOH/Na2O
What do we mean by appreciable?o A 50-mL solution of 0.25 M Acetic
acid.o A 50-mL solution of 0.25 M Sodium
Acetateo A solution containing 0.125 M Acetic
acid and 0.125 M AcetateHA / A–
+ acid+ base
** If we take the ratio of base to acid or acid to base, it should be within 10% of each other
Buffers contain appreciable amounts of a weak acid and its conjugate base
Buffers work because there are weak acids and weak bases present to counter-act small
amounts of acid/bases.
The effectivity (and pH) of the buffer is dependent on the ratio between the weak
acid and its conjugate base
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Ka =H3O
+[ ] A
−[ ]
HA[ ]
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−log Ka( ) = −logH3O
+[ ] A
−[ ]
HA[ ]
⎛
⎝ ⎜ ⎜
⎞
⎠ ⎟ ⎟
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−logKa = −log H3O+
[ ] − logA−
[ ]
HA[ ]
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pKa = pH − logA−
[ ]
HA[ ]
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pH = pKa+ logA−
[ ]
HA[ ]
Henderson-Hasselbach Equation:
The effectivity (and pH) of the buffer is dependent on the ratio between the weak
acid and its conjugate base
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pH = pKa+ logA−
[ ]
HA[ ]
o A 0.25 M Acetic acid buffer with pH 4.74
o A 0.25 M Acetic acid buffer with pH 5.10
o A 0.25 M Acetic acid buffer with pH 4.40
** If we take the ratio of base to acid or acid to base, it should be within 10% of each other
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0.1<A−
[ ]
HA[ ]<10
1 pKapHBUFFER RANGE
The effectivity (and pH) of the buffer is dependent on the ratio between the weak
acid and its conjugate base
o A buffer for pH 10.00
o A buffer for pH 4.00
o A buffer for pH 7.00
The closer the pH of the buffer to the pKa the better
The effectivity (and pH) of the buffer is also dependent on the total amount of weak acid
and conjugate base.
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pH = pKa+ logA−
[ ]
HA[ ]
o A 1.00 M Acetic acid buffer with pH 4.74
o A 0.30 M Acetic acid buffer with pH 4.74
o A 0.10 M Acetic acid buffer with pH 4.74
o A 0.030 M Acetic acid buffer with pH 4.74
Buffer Capacity is the measure of the ability of a buffer to resist pH Changes
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pH = pKa+ logA−
[ ]
HA[ ]
1 pKapH
GOOD BUFFER RANGE+ appreciable amounts (High Concentration)…
Buffers work through a phenomenon known as the common ion effect
What is the common ion effect?This effect occurs when a reactant
containing a given ion is added to an equilibrium mixture that already contains that ion, and the position of equilibrium shifts away from forming more of it
Preparation of Buffers
1. Choose the conjugate acid-base pair2. Calculate the ratio of the buffer
component concentrations3. Determine the buffer concentration4. Mix the solution and adjust pH
Buffers can be prepared using a weak acid and the salt of its conjugate base (or a weak base
and the salt of its conjugate acid).
Example 1Preparing a pH 10.00 carbonate buffer. How many grams of Na2CO3 must one add to 1.5 L of freshly prepared 0.20 M NaHCO3 to make the buffer? Ka of HCO3
- is 4.7 x 10-11
Example 2Prepare a 50-mL of 0.12 M Acetic acid buffer with equal concentrations of acetic acid and acetate from 3.00 M acetic acid stock solution and sodium acetate salt
Buffers can also be prepared by using a weak acid (or weak base) then add a strong
base (or acid) to desired pH.
5.0 g of CH3COONa is dissolved in 100. mL of water. How many mL of 0.50 MHCl should be added to form a buffer with pH 4.90? Will diluting the final mixture to 500 mL affect the pH of the buffer? What is affected by dilution?
Acid-Base TitrationsChapter 10
In an acid-base reaction, the key parameter that changes in the system is pH.
Example:
A 40.00 mL sample of 0.1000 M HCl solution was titrated with 0.1000 M NaOH solution.
Calculate the pH when the following volume of the NaOH is added.
a) 0.00 b) 10.00c) 20.00d) 30.00e) 35.00
f) 39.00g) 40.00h) 41.00i) 45.00j) 50.00
In an acid-base reaction, pH is monitored through colored indicators.
Acid-base indicators are usually weak acids (HIn) which have different color than its conjugate
base (In-).
acidic
basic
change occurs
over ~2 pH units
Acid-base indicators are usually weak acids (HIn) which have different color than its conjugate
base (In-).
40 mL of 0.1000 M HCl
TITRATION OF A STRONG ACID BY A STRONG BASE
TITRATION OF A WEAK ACID BY A STRONG BASE
REGION 1: Before addition (weak acid)
REGION 2: Before Equivalence point (buffer)
REGION 3: At Equivalence point (weak base)
REGION 4: After Equivalence point (strong base)
TITRATION OF A WEAK ACID BY A STRONG BASE
EXAMPLE in Book: Titration of 50.00 mL of 0.0200 M MES (pKa = 6.27) with 0.1000 M NaOH
REGION 1: Before addition (weak acid)
REGION 2: Before Equivalence point (buffer)
REGION 3: At Equivalence point (weak base)
REGION 4: After Equivalence point (strong base)
TITRATION OF A WEAK BASE BY A STRONG ACID
REGION 1: Before addition (weak base)
REGION 2: Before Equivalence point (buffer)
REGION 3: At Equivalence point (weak acid)
REGION 4: After Equivalence point (strong acid)
TITRATION OF A WEAK BASE BY A STRONG ACID
Example in Book: Titration of 25.00 mL of 0.08364 M pyridine (Kb = 1.6 x 10-9) with 0.1067 M HCl.
REGION 1: Before addition (weak base)
REGION 2: Before Equivalence point (buffer)
REGION 3: At Equivalence point (weak acid)
REGION 4: After Equivalence point (strong acid)
TITRATION OF A POLYPROTIC SYSTEMS
TITRATION OF A POLYPROTIC SYSTEMS