Strong or weak? Acid base equilibrium. Unit IV Investigation IV-X ChemCatalyst You have a beaker containing 0.10 M HCl, hydrochloric acid. When you test.

Strong or weak? Acid base equilibrium

Unit IV • Investigation IV-X

ChemCatalyst

You have a beaker containing 0.10 M HCl, hydrochloric

acid. When you test the conductivity of this solution, the

light bulb shines brightly. How do you explain this

observation?

You have a second beaker containing 0.10 M CH3COOH,

acetic acid. When you test the conductivity of this solution,

the light bulb shines, but very dimly. How do you explain

this observation?

Purpose: This activity allows you to compare acid solutions from a molecular point of view.

Unit IV • Investigation IV-X

Making sense

Some acids on the Handout are labeled “strong” and others labeled “weak”. What is the difference between them?

Unit IV • Investigation IV-X

Acids dissociate into ions in solution.

Unit IV • Investigation IV-X

21

Flask 1: 0.010 M HCl hydrochloric acid -strong

Flask 2: 0.002 M HCl hydrochloric acid - strong

pH = 2.7pH = 2

Some acids do not dissociate completely in solution.

Strong and weak acidsAcids that dissociate completely into ions are called strong acids

HCl, HNO3, HBr

Acids that do not dissociate completely in solution are called weak acids

HF, CH3COOH

The pH of an acid solution is determined by:

1)the molarity of the solution

2)the identity of the acid in solution

Unit IV • Investigation IV-X

Unit IV • Investigation IV-X

Check-In

A solution of hydrocyanic acid has a molarity of 0.010 M and a pH of 5.7. Do you think it is a strong or a weak acid? Explain your thinking.

What factors affect the strength of an acid?

Not all acids dissociate to the same extent in solution. Some acids ionize completely, while others ionize partially.

Both molarity and the identity of the dissolved acid affect the pH of a solution.

Strong acids are acids that dissociate completely in solution. They have higher H+ concentrations as a result.

Weak acids are acids that do not dissociate completely in solution. They have lower H+ concentrations as a result.

Unit IV • Investigation IV-X

Some background

• Acid = proton donor• Base = proton acceptor

Conjugate pair

CH3OOH +H2O H3O+ + CH3OO-

conjugate pair

Water does the same thing!

• Lets look at the “equation” for water ionizing

• H2O + H2O H3O+ + OH-

• [H3O+][OH-]/[H2O]2 =K = constant value (55.6 M no mater what), so lets focus on the stuff that changes

• [H3O+][OH-] = Kw = 1.0 * 10-14 (have you seen this before?)

We already knew…

• Neutral means [H3O+]= [OH-]= 1.0 * 10-7

• Acidic means [H3O+] > [OH-]

• Basic means [H3O+]< [OH-]

Some “p’s”

• pH = -log[H+] pOH = -log[OH-]

• [H+] = 10-pH [OH-] = 10-pOH

• pKw = -log(1.0 * 10-14) = 14 = pH + pOH at 25O

Strong vs. weak

• HCl H+ + Cl-

• [H+][Cl-]/[HCl] ~ 1.0

• In other words, we have (negligibly) little HCl left after the system has come to equilibrium

• [H+] = [HCl]

Strong vs. weak

CH3COOH +H2O H3O+ + CH3COO-

OR• HC2H3O2 H+ + C2H3O2

-

• Based upon the demonstration, there is lots of CH3COOH left after the system has come to equilibrium

• How do we know?

• So [H+] ≠ [HC2H3O2]

Just how weak?

• A 1.0 M solution of acetic acid has a pH of 2.37. What percentage of the acetic acid is ionized in the solution?

• 10 -2.37 = [H+] =0.0043 M

• ([H+]/ [HC2H3O2])*100 = % ionization =

(0.0043/1.0)*100 = 0.43%