Revised . AMB 7-2005 Introduction: Even in quite dilute aqueous solutions, acetic acid is very slightly ionized (it would approach 99% ionization only as the concentration approaches 0.0 M): HC2H3O2(aq) + H2O(l) Ka = 1.8 x 10-5. H3O+ + C2H3O2-
In general, if the acid is not extremely weak, the pH of a solution of a weak acid is governed by the concentration of the acid and Ka. Under similar conditions, the pH of a solution of a weak base is determined by the concentration of the weak base and Kb. In solutions containing both a weak acid and a strong acid, both acids play a role in determining the pH of the solution; however, if the concentration of the strong acid is relatively large, it will inhibit the dissociation of the weak acid.(the common-ion effect). The pH of this solution would then be calculated as if the weak acid were not present! (Likewise, in a solution containing both a strong base and a weak base, the strong base concentration would be used to calculate pH.) Buffer solutions contain both a weak acid and its conjugate weak base in appreciable concentrations. Within limits, these solutions tend to resist changes in pH upon addition of either H3O+ or OH- (because these species are largely consumed by the acidic and basic components of the buffer mixture). In buffer systems like NaC2H3O2-HC2H3O2 mixtures, the principal source of the acetic acid molecule is from the acid; the principal source of the acetate ion is from the salt. Therefore, the [H3O+] is determined by the salt/acid (or equivalently, the base/acid) mole ratio. For a conjugate acid/base pair: pH = pKa + log10 [nconjugate base] [nacid]
On the other hand, the salt, sodium acetate (NaC2H3O2), is essentially 100% dissociated into the constituent hydrated ions in a dilute aqueous solution. Sodium acetate solutions are basic because the acetate ion (the conjugate base of acetic acid) behaves as a proton acceptor with respect to water: C2H3O2- + H2O (l) Kb = 5.6 x 10-10. OH- + HC2H3O2(aq)
Consequently, in solutions of acetic acid alone, the molar concentration of the HC2H3O2 is much larger than the C2H3O2- ion concentration, but in solutions of sodium acetate alone, the reverse is true. Neither solute alone can provide "comparable" concentrations in solutions of both the weak acid and its conjugate base; thus neither an acetic acid solution nor a sodium acetate solution is a buffer solution,i.e. a buffer solution, by definition, must contain moderate concentrations of both species!
where n stands for moles (or millimoles). and the symbol, p, to indicate "-log10 of". For acetic acid at 25°C, pKa = 4.74. This experiment deals with the pH of a solution of acetic acid alone and with a solution of sodium acetate alone (Part A); of "direct" buffer solutions prepared by mixing solutions of acetic acid with solutions of sodium acetate (Part B); of "indirect" buffers prepared by partially "neutralizing" an acetic acid solution with NaOH (Part C), and with the comparison of the pH-changes observed when strong acid (HCl) solutions are added to equal volumes of buffered and unbuffered solutions.
Transfer 10-20 mL of these solutions to separate clean 50 mL beakers; measure and record the pH of each solution - these data will be used to compare the observed pH with that calculated using the Ka and Kb values provided. Part B.. pH Measurements on Buffer Mixtures of Acetic Acid and Sodium Acetate. Using the 0.25 M solutions prepared in Part A, prepare in clean 50 mL beakers the following ("direct") buffer solutions: Beaker # mL of 0.250 M HC2H3O2 (aq) B1 B2 16.0 12.0 10.0 6.0 2.0 mL of 0.250 M NaC2H3O2 (aq) 4.0 8.0 10.0 14.0 18.0
Experimental (work in pairs).
You will use the same pH probes as in the previous experiment. Please remember that the probe must be thoroughly rinsed and recalibrated using the pH =7 buffer periodically. In general, volume measurement will be made with 25 or 50 mL...