Expt 4-Ionic Equilibria

Experiment 4-Ionic Equilibria

Crystle T. CotingtingReinavelle Jeunesse D. Mateo

Objectives• Distinguish acidic from basic compounds• Distinguish a strong acid/base from a

weak acid/base• Determine experimentally the ionization

constant of a weak electrolyte• Determine experimentally the solubility

product constant of a slightly soluble substance

Review: Acids and Bases

• Arrhenius– Acids: substances that contain hydrogen and

dissociate in water to yield hydronium ions – Bases: substances that contain hydroxyl groups

and dissociate in water to yield hydroxide ions• Bronsted-Lowry

- expanded the definition of a base - Base: specie with an electron pair available to

accept a transferred proton.

Review: Acids and Bases• Lewis

- Acids: electron pair acceptors - Bases: electron pair donators

Review: Strength of Acids & Bases• Acids & Bases –

electrolytes • Strong acids & bases –

almost completely dissociate or ionize in water

• Weak acids & bases –dissociate only slightly

Ionic Equilibrium• observed when ionization takes place in a

solvent medium.

• equilibrium constant, K, for the ionization of a weak acid or base.

• tells us how far to the right the reaction will proceed to reach the equilibrium

Ionization or Dissociation Constant

Ionization Constant

• In an acid, HA(aq)+ H2O(l) A-

(aq)+ H3O+(aq)

it has the general form

Kacid = [H3O+] [A-]

[HA]

Solubility Product Constant

• Ionic solid + pure water dissolution at a relatively rapid initial rate

• As the concentration of dissolved ions increases, so does the rate of formation.

• Soon, rateformation = ratedissolution (state of equilibrium and saturation of solution)


• At equilibrium, no net dissolution of the solid. If more of the solid is added to the mixture of solid and solution at equilibrium, none will dissolve.

• If solid is removed, the concentration of ions in solution will stay the same.


• Presence of solid - no effect on the equilibrium concentrations of ions in the saturated solution

• Therefore, the equilibrium constant expression for equation does not include a term referring to the solid.

• Solubility product constant is defined as such:

• Ksp = [Am+]n[Bn-]m

Solubility Product Constant• A measure of how far to the right the

dissolution proceeds at equilibrium• Important in explaining phenomena like

solubility and precipitation of compounds

ExperimentalMethodology

Part 1. Strong and Weak Electrolytes

• Put pH paper to solutions A-G

• Determine pH of each solution

• Classify if strong or weak acid or base or neutral

Part 2. Ionization Constant of Acetic Acid

1.1 ml 1M acetic acid

1 drop methyl violet

1.0 ml distilled water

1 drop methyl violet

0.01M HCl from syringe

Match color

Part 3. Solubility Product Constant of Benzoic Acid, C6H5COOH

50 ml distilled water

40°C

Pinch of benzoic acid crystals

2 drops phenolph-thalein

0.01M NaOH titration to light pink end point

Results and Discussion

Part 1. Strong and Weak Electrolytes (Result)

Observed pH for Various Solutions

Solution Observed pH

A 1

B 4

C 13

D 7

E 8

F 6

G 6

Part 1. Strong and Weak Electrolytes (Discussion)

Classification System Devised for Strong or Weak Acid and Strong or Weak Base

pH Classification

0-2.5 Strong acid

2.5-6 Weak acid

7 Neutral

8-11 Weak base

12-14 Strong acid


Final Classification of SolutionsSolution Observed pH Classification

A 1 Strong acid

B 4 Weak acid

C 13 Strong base

D 7 Neutral

E 8 Weak base; basic salt

F 6 Weak acid; acidic salt

G 6 Weak acid; acidic salt


• Ka of solution A > Ka of solutions B, F, and G.

Solution Observed pH Classification

A 1 Strong acid Ionization complete

B 4 Weak acid Slight ionization

C 13 Strong base Ionization complete

D 7 Neutral

E 8 Weak base; basic salt Slight ionization

F 6 Weak acid; acidic salt Slight ionization

G 6 Weak acid; acidic salt Slight ionization

Part 2. Ionization Constant of Acetic Acid (Result)• Volume of HCl added (to match light purple

color of acetic acid) - 0.6mL of HCl

Acetic Acid

Part 2. Ionization Constant of Weak Acid (Discussion)Methyl violet• a pH indicator to test pH ranging from 0 to 1.6.

- acid end (pH below 0.0): yellow

- alkaline end (above pH 1.6): bluish violet

Part 2. Ionization Constant of Weak Acid (Discussion)

• Addition of Methyl Violet– 1.1mL 1M acetic acid: light purple color

solution has a pH about or above 1.6

acetic acid – weak acid – About 0.6mL of 0.01M HCl was added to 1.0mL of

distilled water until the solution turned the same color as acetic acid

2 solutions with same pH indicator are the same color about the same pH

pH of acetic acid = pH of HCl solution

Since pH = -log [H3O+],

[H3O+] in HCl solution = [H3O+] in the acetic acid

(HOAc) solution

[H3O+] in HCl solution = molarity of the HCl because HCl is a known strong acid and strong acids dissociate almost completely into their respective ions


Calculation for Molarity of HCl

Note: Can’t immediately assume to be 0.01M because it has been diluted by the 1.0 mL distilled water

• C1V1 = C2V2

– C1 = 0.01 M original concentration of HCl

– V1 = 0.6 mL volume of HCl added

– V2 = 1.6 mL 1.0mL H2O + 0.6mL HCl

• (0.01M)(0.6mL) = (C2)(1.6mL)

• C2 = (0.01M)(0.6mL)

(1.6mL)

• C2 = 0.00375 M diluted concentration of HCl


• C2 = 0.00375 M

• C2 = [HCl] = 0.00375 M

• [HCl] = [H3O+] in HCl solution = 0.00375 M– Because it is a strong acid

• [H3O+] in HCl solution = [H3O+] in the acetic acid

(HOAc) solution – Because same color with same pH indicator

• Therefore, [H3O+] in the acetic acid solution = 0.00375M


Part 2. Ionization Constant of Acetic Acid (Discussion)

Chemical equation for the ionization of acetic acid (HOAc)

HOAc(aq) + H2O(l) OAc-(aq) + H3O+

(aq)

Part 2. Ionization Constant of Acetic Acid (Discussion) Calculation of [OAc-] and [HOAc].

HOAc(aq)+ H2O(l) OAc-(aq)+ H3O+

(aq)

Initial 1M - 0 0

Change -x - +x +x

Equilibrium 1M-x - x x

At equilibrium, x = [H3O+] [H3O+] = 0.00375 M.

Since x = [OAc-] = [H3O+] [OAc-] = 0.00375 M.At equilibrium,

[HOAc] = 1M – x = 1M – 0.00375M = 0.99625 M

Part 2. Ionization Constant of Acetic Acid (Discussion)

Summary of ValuesVolume of HCl used 0.6 mL

[H3O+] in HCl solution 0.00375 M

H3O+] in HOAc solution 0.00375 M

Equilibrium Concentration of HOAc solution

0.99625 M

From here, we can solve for Ka. Ka = [OAc-] [H3O+] [HOAc] = (0.00375 M) (0.00375M) (0.99625 M) = 1.4115 x 10-5

The theoretical Ka for acetic acid at 25ºC is 1.77x10-5

Part 2. Ionization Constant of Acetic Acid (Conclusion)

• Calculated Ka value of acetic acid is 1.41x10-5

- dissociates only slightly weak acid

Part 3. Solubility Product Constant of Benzoic Acid (Results)

• 19.7 mL of the titrant, 0.01M NaOH was needed to reach the light pink end point of titration

Part 3. Solubility Product Constant of Benzoic Acid (Discussion)

• Calculate concentration of H3O+ using the following equation MacidVacid=MbaseVbase.

MacidVacid=MbaseVbase

Mbase = 0.01M

Vbase= 19.7mL

Vacid = 10mL

(Macid)(10mL) = (0.01M)(19.7mL)

Macid = 0.0197M

• At end point, solution being titrated is neutralized by the titrant.

• Therefore,

[OH-] of titrant = [H3O+ ] of acid being titrated

• In this experiment,

[OH-] NaOH = [H3O+] benzoic acid = 0.0197 M


• Get value of [C6H5COO-] using the initial-change-equilibrium table.

[C6H5COO-] = x

[H3O+] = x

[C6H5COO-] = [H3O+]

= 0.0197 M too


• Calculating the solubility product constant

Ka = [C6H5COO- ][ H3O+]

[C6H5COOH]

Ka[C6H5COOH] = [C6H5COO-] [H3O+]

Ksp = [C6H5COO-][H3O+]

= (0.0197M) (0.0197M)

= 0.00038809

~ 3.9 x 10-4


Calculating the solubility of benzoic acid

Ksp = [C6H5COO-] [H3O+] *S – solubility

0.00038809 = S x S

0.00038809 = S2

S = 0.0197 M

Theoretical value (25°C): 0.02785 M


• Ksp value for benzoic acid is 0.003889. Similarly, since Ksp is a measure of how far to the right the dissolution proceeds at equilibrium, a compound with a higher Ksp value is more soluble than this system.

Part 3. Solubility Product Constant of Benzoic Acid (Conclusion)

Recommendations

Recommendations • The group recommends the use of a pH meter

instead of pH paper to determine the pH of the seven solutions. Or, if a pH meter is not available we recommend the use of other indicators, so as to more accurately estimate the pH. Comparing our data with other groups, we noticed that the pH we got were different from the pH they observed, probably because comparing the colors in the pH paper is very subjective to the user as well as the lighting of the room.

• Also, in determining the ionization constant of acetic acid, titrating with a base would prove more effective than matching colors, because comparing colors again is quite subjective to the one performing the experiment.

Recommendations

The End.

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Expt 4-Ionic Equilibria