Strong or weak? Acid base equilibrium

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? Unit IV • Investigation IV-X 2

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. 2 1 Flask 1: 0.010 M HCl hydrochloric acid -strong Flask 2: 0.002 M HCl hydrochloric acid - strong pH = 2.7 pH = 2 Unit IV • Investigation IV-X

Some acids do not dissociate completely in solution.

Strong and weak acids Acids 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: the molarity of the solution the identity of the acid in solution 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. Unit IV • Investigation IV-X

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%