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Acid-Base Titrations
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Introduction to Acids and Bases
Chapter 8
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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.
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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
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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
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Acids and Bases undergo neutralization reactions.
NaOH + HCl NaCl + H2O
Na+ + OH– + H+ + Cl– Na+ + Cl– + H2O H+ + OH– H2O
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Acids and Base strengths are defined by how much H3O+ or OH- is produced by a given concentration
Strong Acid Weak Acid
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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
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Weak acids and bases are in equilibrium with their original species.
CH3COOH + H2O H3O+ +CH3COO–
€
Kc =H3O
+[ ] CH3COO
−[ ]
H2O[ ] CH3COOH[ ]Kc << 1
€
Kc H2O[ ] =H3O
+[ ] CH3COO
−[ ]
CH3COOH[ ]
€
Ka =H3O
+[ ] CH3COO
−[ ]
CH3COOH[ ]
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Weak acids and bases are in equilibrium with their original species.
NH3 + H2O NH4+
+OH–
€
Kc =NH4
+[ ] OH
−[ ]
H2O[ ] NH3[ ]Kc << 1
€
Kc H2O[ ] =NH4
+[ ] OH
−[ ]
NH3[ ]
€
Kb =NH4
+[ ] OH
−[ ]
NH3[ ]
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Ka and Kb tells us something about the relative acid/base strengths
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Ka and Kb tells us something about the relative acid/base strengths
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Water can act both as an acid and a base (AUTOIONIZATION)
€
Kc =H3O
+[ ] OH
−[ ]
H2O[ ]2
€
Kc H2O[ ]2
= H3O+
[ ] OH−
[ ]
€
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
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The p-scale conveniently handles a wide range of concentrations
]log[ 3 OHpH
]log[ OHpOH
KwpKw log KapKa log
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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.
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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?
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Complete the following table:
Parameter
An NaOH solution
A 0.30 M HNO3
pH 9.52
pOH
[H3O+]
[OH-]
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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
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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
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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
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BuffersChapter 9
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Buffers solutions are solutions that resist changes in pH
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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
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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
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Buffers work because there are weak acids and weak bases present to counter-act small
amounts of acid/bases.
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The effectivity (and pH) of the buffer is dependent on the ratio between the weak
acid and its conjugate base
€
Ka =H3O
+[ ] A
−[ ]
HA[ ]
€
−log Ka( ) = −logH3O
+[ ] A
−[ ]
HA[ ]
⎛
⎝ ⎜ ⎜
⎞
⎠ ⎟ ⎟
€
−logKa = −log H3O+
[ ] − logA−
[ ]
HA[ ]
€
pKa = pH − logA−
[ ]
HA[ ]
€
pH = pKa+ logA−
[ ]
HA[ ]
Henderson-Hasselbach Equation:
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The effectivity (and pH) of the buffer is dependent on the ratio between the weak
acid and its conjugate base
€
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
€
0.1<A−
[ ]
HA[ ]<10
1 pKapHBUFFER RANGE
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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
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The effectivity (and pH) of the buffer is also dependent on the total amount of weak acid
and conjugate base.
€
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
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Buffer Capacity is the measure of the ability of a buffer to resist pH Changes
€
pH = pKa+ logA−
[ ]
HA[ ]
1 pKapH
GOOD BUFFER RANGE+ appreciable amounts (High Concentration)…
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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
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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
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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
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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?
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Acid-Base TitrationsChapter 10
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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
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In an acid-base reaction, pH is monitored through colored indicators.
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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
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Acid-base indicators are usually weak acids (HIn) which have different color than its conjugate
base (In-).
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40 mL of 0.1000 M HCl
TITRATION OF A STRONG ACID BY A STRONG BASE
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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)
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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)
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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)
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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)
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TITRATION OF A POLYPROTIC SYSTEMS
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TITRATION OF A POLYPROTIC SYSTEMS