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Lab 24 - Hydrolysis. A salt formed between a strong acid and a weak base is an acid salt . Ammonia is a weak base, and its salt with any strong acid gives a solution with a pH lower than 7. For example, let us consider the reaction. H 2 O + NH 3 NH 4 + + OH -.
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Lab 24 - Hydrolysis A salt formed between a strong acid and a weak base is an acid salt. Ammonia is a weak base, and its salt with any strong acid gives a solution with a pH lower than 7. For example, let us consider the reaction • H2O + NH3 NH4+ + OH-
In the solution, the NH3+ion reacts with water (called hydrolysis) according to the equation: NH3++ H2O NH4+ OH-. • The acidity constant can be derived from Kw and Kb. • Ammonia Kb=1.75e-5 • KwKa = ---------------- Kb = 1.00e-14 / 1.75e-5 = 5.7e-10.
What is the pH 0f a .1M NaCLO solution if Ka for HCLO 1s 3.0e-8. • step 1 write the net ionic formula • Na + + CLO- + H2O HCLO + Na + OH- • step 2 determine initial, change and equilibrium [M] of reactants and products
What is the pH 0f a .1M NaCLO solution if Ka for HCLO 1s 3.0e-8. • step 3 Find the Kb value KwKb= ---------------- ------ Ka 1.0e-143.3e-7.= ---------------- ------3.0e-8.
What is the pH 0f a .1M NaCLO solution if Ka for HCLO 1s 3.0e-8. • step 3 Find the M of [OH} [HCLO] [OH-]Kb= ---------------- ------ [CLO] [X2] 3.3e-7= ---------------- [0.1] = 1.8 e -4
What is the pH 0f a .1M NaCLO solution if Ka for HCLO 1s 3.0e-8. • step 4 Convert [OH] to pOH -log 1.8 e -4 = 3.74 step 5 Convert pOHto pH 14 -3.74 = 10.26
A. Hydrolysis of salts 1. Add 5ml of Di water into 6 test tubes 2. Add 3 drops of indicator (as listed on page 270) to each test tube 3. Using the indicators chart on page 262 determine the pH of the water 4. Repeat steps 1-3 using boiled water 5. Repeat steps 1-3 using .1M solutions of the salts listed on page 269
A. Data 1. Fill out chart on pg 269 based on the hydrolysis behavior on page 260 2. Determine (H +) for each pH 3. Determine (OH -) for each pH 4. Using data from pg 269fill out chart on page 271
calculations • [H+] = 10-pH ( anti log of –pH) • [OH-] = Kw / [H+] • [OH-] = [1.0 x 10-14] / [H+] • Ka or Kb = [M] Products over [M] reactants • Omit water and assume both products have the same molarity
Lab 25 – dissociation of a weak acid General Theory According to the Brønsted-Lowry acid-base theory, the strength of an acid is related to its ability to donate protons. All acid-base reactions are then competitions between bases of various strengths for these protons. For example, the strong acid HCl reacts with water according to Equation [1]: HCl(aq) + H2O(l)H3O+(aq) + Cl-(aq) General
Lab 25 • This is a strong acid and is completely dissociated (in other words, 100 percent dissociated) in dilute aqueous solution. Consequently, the [H3O+] concentration of a 0.1 MHCl solution is 0.1 M. Thus HCl is a stronger acid than water and completely donates a proton to water to form H3O+.
This is a strong acid and is completely dissociated (in other words, 100 percent dissociated) in dilute aqueous solution. Consequently, the [H3O+] concentration of a 0.1 MHCl solution is 0.1 M. Thus HCl is a stronger acid than water and completely donates a proton to water to form H3O+.
By contrast, acetic acid, HC2H3O2 (abbreviated HOAc), is a weak acid and is only slightly dissociated, as shown in Equation [2]: • Its acid dissociation constant, as shown by Equation [3], is therefore small • Ka= = 1.8 x 10-5
titrate the weak acid • If we titrate the weak acid HA with a base, there will be a point in the titration at which the number of equivalents of base is just one-half the number of equivalents of acid present (at ½ Ve). • pH = pKa@ ½ [H+] • By titrating a weak acid with a strong base and recording the pH versus the volume of base added, one can determine the dissociation constant, Ka, of the weak acid
. C. Determination of pKa of Unknown • Pipeta 25.00 mL aliquot of your unknown acid solution into a 250-mL beaker and carefully immerse the previously rinsed electrode into this solution. • Measure the pH of this solution. Record the pH in your notebook. Begin your titration by adding 1 mL of your standardized base from a buret and record the volume of titrant and pH. • Repeat with successive additions of 1 mL of base until you approach the end point. • equivalence point--where the graph is steepest • See pg 277 • Add 0.1-mL increments of base and record the pH and milliliters of NaOH added untill the pH no longer changes • plot graph to determine ½ equivalence point-
Formulas • pH = pKa @ ½ equivalence point- • ½ equivalence point = 4.3 • pKa = 4.3 • Ka = antilog - 4.3 = 5 e-5
Due • Pg 264,265 • Questions 1-3