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–

€

Kc =H3O

+[ ] CH3COO

−[ ]

H2O[ ] CH3COOH[ ]Kc << 1

€

Kc H2O[ ] =H3O

+[ ] CH3COO

−[ ]

CH3COOH[ ]

€

Ka =H3O

+[ ] CH3COO

−[ ]

CH3COOH[ ]

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[ ]

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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)

€

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

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

€

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:

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

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.

€

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

€

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