Principles of Neutralization Titration: Determining Acids, Bases and pH of

Buffer Solutions

What are acids and bases?

• An acid is a substance that increases the concentration of H3O+ (hydronium ion) in aqueous solution.

• A base: is a substance that decreases the concentration of H3O+ in aqueous solution.

A decrease in [H3O+] requires an increase in [OH-], so we can define the base as a substance that increase the concentration of OH- in aqueous solution.

- H+ is called a proton, because a proton is all that remains when a hydrogen atom loses its electron.

- Hydronium ion (H3O+) is a combination of H+ with H2O.

Bronsted and Lowry definition of acids and basesAn acid is a proton donorA base is a proton acceptor.

ExampleHCl is an acid because it donates a proton to H2O to form H3O+

HCl(l) + H2O H3O+

(aq) + Cl-(aq)

Bronsted and Lowry definition can be extended to nonaqueous solvents and to the gas phaseExample HCl (g) + NH3(g) NH4

+Cl-

Hydrochloric acid ammonia ammonium chloride (acid) (base) (salt)Salts:- Any ionic solid such as ammonium chloride is called salt. It can be thought of as the product of an acid base reaction.- Most salts are strong electrolytes, i.e. they dissociate almost completely into their component ions when dissolved in water

NH4+Cl- NH4

+ + Cl-

H20

Conjugate acids and bases:The products of a reaction between an acid and a base are also acid and base

CH3COOH + CH3NH2 CH3COO- + CH3NH3

Acetic acid Methylamine acetate ion methyl ammonium ion ( acid) (base) (base) (acid)

Conjugate pair

Conjugate pair

Acetate is a base because it can accept a proton to make acetic acid.methyl ammonium ion is an acid because it can donate a proton and become Methylamine base.

acetic acid and acetate ion are said to be a conjugate acid-base pair.Methylamine and methyl ammonium ion are also said to be a conjugate acid-base pair.

Conjugate acids and bases are related to each other by the gain or lose of one H+.

Acid Dissociation Constants

Acid Dissociation Constants (continue)

What solutions and indicators are used in neutralization titrations?

• Neutralization titration depend on a chemical reaction between the analyte and a standard reagent .

• The point of chemical equivalence is indicated by a chemical indicator or an instrumental measurement.

• The standard solutions employed in neutralization titrations are strong acids or strong bases (because they react more completely with an analyte than their weaker counterparts do and thus yield sharper endpoints).

• standard solutions of acids or bases are prepared by diluting a concentrated solution. e.g. acids, HCl, HClO4, H2SO4) bases (NaOH, KOH, BaOH), remember that those solutions need standardization.

• Week acids and bases are never used as standard reagents because they react incompletely

Titration• A standard solution (standard titrant) is a reagent of known concentration

that is used to carry out a titrimetric analysis

• Titration: is a procedure performed by adding a standard solution from a buret or other liquid-dispensing devices to a solution of the analyte until the reaction between the two is judged complete.

• Equivalence point in titration is reached when the amount of added titrant is chemically equivalent to the amount of analyte in the sample

• E.g. AgNO3 + Cl- AgCl mol mol• (Equivalent points can’t be determined experimentally)

• End point is the point in a titration when physical change occurs that is associated with the condition of chemical equivalence.

• Indicator: a chemical compound that change it’s color or other physical property at or near the equivalence point.

Acid / Base Indicator• Acid base indicator is a week organic acid or a weak

organic base whose un-dissociated form differs in color from its conjugate form.

• It could be natural or synthetic compound which display colors that depends on the pH of the solution in which they are dissolved.

• Acid type indicator HIn HIn + H2O In- + H3O+ Ka= [H3O+] =Ka

acid color Base color

• Base type indicator In-

In- + H2O HIn + OH-

Base color acid color

[H3O][In-]

[HIn]

[HIn]

[In-]

Structures of some of the acid - base indicator

e.g. Bromocresol green

Yellow BluepH < 3.8 pH > 5.4

• Acid type indicator HIn HIn + H2O In- + H3O+ Ka= [H3O+] =Ka

The hydronium ion (H3O+) concentration determines the ratio of the acid to the conjugate base form of the indicator and thus determine the color developed by the solution.

The indicator HIn exhibit its pure acid color when [HIn]/[In-] ≥ 10

The indicator HIn exhibit its pure base color when [HIn]/[In-] ≤ 0.10.

The color appears to be intermediate for ratios between these two values.

The full acid color [H3O+] = 10Ka

The full base color [H3O+] = 0.1Ka

[HIn]

[In-]

[H3O][In-]

[HIn]

Acid / Base Indicator• To obtain the indicator pH range we take the negative logarithm of the two

expressions:

pH (acid color) = - log(10Ka) = pKa +1

pH (base color) = - log(0.1Ka) = pKa – 1

Indicator pH range = pKa ± 1

(i.e. an indicator with an acid dissociation constant 1.0 x 10-5 (pKa =5) typically shows a complete color change when the pH of the solution it is dissolved in changes from 4 to 6).

The following tables show some of the common acid/base indicators. The transition range ( i.e. range for the indicator to go from acidic to basic color) ranges from 1.1 to 2.2 ( i.e. pH = pKa ± 1.1 or pH = pKa ± 2.2)

Calculating pH in Titrations of strong acids and strong bases

• The hydronium ion in an aqueous solution of a strong acid has two sources:

• (1) the reaction of the acid with water• (2) the dissociation of water it self.

Treating strong acid with strong base

• Here we will be interested in calculating the hypothetical titration curves of pH versus volume of titrant.

• Three types of calculation must be done to construct the hypothetical curve for treating a solution of strong acid with a strong base:

1- pre-equivalence ( we compute the conc. Of the acid from its starting conc. And the amount of the base that has been added)

2- equivalence (H3O+ and OH- are present at equal concentrations, [H3O+] is derived directly from the ion product constant of water)

3- post-equivalence (conc. of excess base is computed. [OH-] is converted to pH by Kw = [H3O+][OH-] pKw=pH + pOH

Example:Do the calculations needed to generate the hypothetical titration curve

for the titration of 50 mL of 0.05M HCl with 0.1 M NaOH.

Titrating strong base with strong acid

• Example:

Calculate the pH during the titration of 50 mL of 0.05 M NaOH with 0.1 M HCl after the addition of the following volume of reagent: (a) 24.5 mL, (b) 25 mL, (c ) 25.5 mL.

Buffer solutions• Buffer solution is a solution of conjugate

acid/base pair that resists changes in pH as a result of dilution or the addition of small amount of strong acid or base.

• It can be prepared either by mixing weak acid with its conjugate base (salt), or by adding small amount of strong base to weak acid or by adding small amount of strong acid to weak base.

• Buffers are used in experiments or applications whenever the pH of a solution is needed to be maintained constant or at a predetermined level.

Calculating the pH of Buffer solutions HA + H2O H3O+ + A- Ka=[H3O+][A-]/[HA]

A- + H2O OH- + HA Kb= [OH-][HA]/[A-]

[H3O+] = Ka. [HA]/[A-]Or

pH = pKa + log ([A-]/[HA])

Example: What is the pH of a solution that is 0.4 M in formic acid and

1.0 M of sodium formate (Ka, HCOOH = 1.8 x 10-4)

Example 2

Calculate the pH of a solution that is 0.2 M in NH3 and 0.3 M in NH4

+Cl- (Ka, NH4+ = 5.7 x 10-10)

What are the unique properties of buffer solution?

• Buffers do not maintain pH at absolutely constant value, but changes in pH are relatively small when amounts of acids or bases are added.

1- Effect of dilution

2- The effect of added acids and bases

3- The buffer capacity, (what is buffer capacity?)

1- The effect of dilution:- The pH of a buffer solution remains essentially

independent of dilution until the concentrations of the species it contains are decreased to the point that the concentrations of OH- and H+ from solution is very close to the concentration of acid and its conjugate base.

2- the effect of added acids and bases:

- Buffer solution resist pH change after addition of small amount of an acid or a base.

Example: Calculate the pH change that takes place when 100 mL portion of 0.05 M NaOH and (b) 0.05 M HCl is added to 400 mL of buffer solution 0.2 M NH3 and 0.3 M NH4

+Cl-.

3- What is the buffer capacity• The ability of the buffer to prevent a significant change in

pH is directly related to the total concentration of the buffering species as well as to the concentration ratio.

• The buffer capacity of a solution is defined as the number of moles of strong acid or base that causes 1 liter of the buffer to undergo a 1 unit change in pH.

• Buffer capacity depends on :1- total concentration of the two buffer components.2- concentration ratios of the two buffer components.

• Buffer capacity falls off moderately rapidly as the concentration ratio of acid to conjugate base departs from unity. For this reason the pKa of an acid chosen for a given application should lie within ±1 unit of the desired pH for the buffer to have a reasonable capacity.

Preparing buffers:By making up a solution of approximately the desired pH and then adjust it by adding acid or conjugate base until the required pH is indicated by pH meter.

Calculating the pH value in weak acid titration• Four distinctly different types of calculations are needed to derive a

titration curve for a weak acid or a weak base.

1- At the beginning, the solution contains only a weak acid or a weak base, and the pH is calculated from the concentration of that solute and its dissociation constant.

2- After various increments of titrant have been added (in quantities up to, but not including equivalent amount), the solution consists of a series of buffers. The pH of each buffer can be calculated from the analytical concentration of conjugate base or acid and the residual concentrations of the weak acid or base.

3- At the equivalence point, the solution contains only the conjugate of weak acid or base being titrated (that is a salt), and the pH is calculated from the concentration of this product.

4- Beyond the equivalence point, the excess of strong acid or base titrant represses the acidic or basic character of the reaction product to such and extent that the pH is governed largely by the concentration of excess titrant

• Example 1: Generate a curve for the titration of 40ml of 0.1 M acetic acid

(Ka= 1.75 x10-5) with 0.1 M sodium hydroxide. Find the pH after adding a) 0.00ml , b) 20.00 ml ,c) 40.00 ml, d) 41.00 ml of titrant. Plot the titration curve.

HPr : hypothetical acid, Ka = 1x 10-5

• Example 2

A 40 ml aliquot of 0.1 M NH3 (Ka, NH4+= 5.7 x10-10) is titrated

with 0.1 M HCl. The reaction is

NH3 + H2O NH4+ + OH-

Calculate the pH after the addition of (a) 0.00, (b) 20.00,

(c ) 40 and (d) 41 mL of acid.