Part 1: Experimental Measurement—Determining a Numerical Value for KwExperimental pH of 0.010 M NaOH = ____________Chemistry 112 Laboratory: Chemistry of Acids & Bases Page 73Name _________________________________________Partner’s Name _________________________________________Grade and Instructor CommentsDATA SHEETS AND CALCULATIONS FOR ACIDS & BASESPart 1: Calculations——Determining a Numerical Value for KwWhat is your measured pH? _______________________Based on the measured pH, what is the hydronium ion concentration? [H3O+] = ______________ MKnowing that in a 0.010 M NaOH solution, [OH-] = 0.010 M, calculate a value for Kw from your experi-mental value of the measured hydronium ion concentration and the known OH- concentration of the 0.010 M NaOH. Kw = [H3O+][OH-] = and pKw = -log Kw = Compare your results with the data taken from the scientific literature: T (˚C) Kw pKw 20 0.68 x 10-14 14.17 25 1.01 x 10-14 14.00Kw, as is the case for all equilibrium constants, varies with temperature. However, at a given temperature, Kw is a constant for an aqueous solution. This means that at 25 ˚C in any aqueous solution, regardless of solute, the value of Kw {=[H3O+][OH-]} is 1.01 x 10-14.What is the pH of pure water at 25 ˚C?Revised: December 2005Chemistry 112 Laboratory: Chemistry of Acids & Bases Page 74Part 2. Determination of Ka for the Ammonium Ion—Experimental MeasurementsData Table for Solutions of Ammonium Ion and AmmoniaEnter the experimental pH values you determine in the lab in the last column. Complete the open areas of the table. * Be very careful to rinse your glass electrode thoroughly with water before an after making this measurement. Solution [NH4Cl], M [NH3], M Solute Type Enter Your (Acid, Base, or Acid + Conj. Base Experimental pH A 0.10 0 Acid B 1.0 0 C 0.050 0.050 D 0.50 0.50 E 0 0.10 Revised: December 2005Chemistry 112 Laboratory: Chemistry of Acids & Bases Page 75Part 2: Calculating Ka for the Ammonium Iona) Write the balanced, net ionic equation for the reaction of ammonium ion with water.b) Write the equilibrium constant expression for Ka for aqueous NH4+c) Enter your experimental information for [NH3] and [NH4+] (from the previous page) into the table below. Use your measured pH values for each solution (A-E) to calculate [H3O+] and enter these values in the table below. Finally, calculate Ka for the ammonium ion and enter the values in the table. Show one representative calculation here. Average calculated Ka value = ______________________ and average pKa = ___________________Revised: December 2005 Solution [H3O+], M [NH3], M [NH4+], M Calculated Ka for NH4+ Calculated pKa for NH4+ A -B C D E Not a required calcu-lationNot a required calcu-lationPart 3. Properties of NH4+/NH3 Buffer Solutions—Experimental Measurements• Your experimental readings from page 70 are entered in the column labeled “Initial pH.”• NOTE: make a pH measurement on pure water before adding acid or base. • Data for the pH after the addition of excess acid or base is entered into the columns marked “pH on adding H+” and “pH adding OH-” as appropriate.• Fill in the boxes marked “∆pH” with calculated numbers.Chemistry 112 Laboratory: Chemistry of Acids & Bases Page 76 Solution [NH4+], M [NH3], M ∆pH on pH on Initial pH pH on ∆pH adding H+ adding H+ adding OH- adding OH- A 0.10 0 -B 1.0 0 C 0.050 0.050 D 0.50 0.50 E* 0 0.10 -Pure water Revised: December 2005Part 3: Properties of NH4+/NH3 Buffer Solutions—Questions and CalculationsEffect of Dilution on the pH of a Buffer: If the solution is diluted more than 10-fold, which solution — 1.0 M NH4Cl (solution B) or 0.50 M NH4Cl + 0.50 M NH3 (solution D) —does the pH change more? (Base your answer on the data in the “Initial pH” column on page 76.)Explain, on the basis of the Ka expression, why dilution has less effect on the pH of a buffer solution than on the pH of a solution containing only the acid as a solute (here NH4+).Effect of Added H3O+ and OH- on a BufferCompare the values of ∆pH (the changes in pH) for solutions C and D (in the table on page 15) with those for solutions of the acid along (A and B) or conjugate base alone (E). a) Which solutions show a buffering action?b) Write balanced chemical equations for reactions that prevent larger changes in pH.Chemistry 112 Laboratory: Chemistry of Acids & Bases Page 77Revised: December 2005Part 4. Titration CurvesThe change in indicator color in an acid-base titration is a signal that the equivalence point is very near. Here you test two indicators that change colors in two different pH ranges.See Chemistry & Chemical Reactivity, page 872, Figure 18.10 for indicator colors.Chemistry 112 Laboratory: Chemistry of Acids & Bases Page 78 Indicator Color in Color in Acidic Solution Basic Solution Bromcresol green Phenolphthalein Revised: December 2005Chemistry 112 Laboratory: Chemistry of Acids & Bases Page 79Titration Results: Option (a)—HCl + NaOHThe volumes of NaOH in the table are suggested values. Enter your actual volumes of NaOH used in the table (second and fifth columns). Enter experimental data in every cell in the table.Deductions from the HCl Titration Curvea) Write a balanced, net ionic equation for the reaction that occurs during the titration.b) How many equivalence points can you detect? Explain the connection between the number of equivalence points and the reaction occurring.c) CLEARLY LABEL on your titration curve the formulas for the species present at: i) before adding NaOH (ii) after 15 mL of NaOH has been added ii) at the equivalence pointd) What is the connection between the indicator colors and the equivalence point? Suggested Actual Measured Indicator Suggested Actual Measured Indicator VNaOH, mL VNaOH, mL pH Color VNaOH, mL VNaOH, mL pH Color 0 22 3 23 6 24 8 24.5 10 25 12 25.5 14 26 16 28 18 30 20 32Be sure to attach to your report form a carefully drawn plot of pH versus volume of base added. Be sure your name appears on the plot.Revised: December 2005Titration Results: Option (b)—H3PO4 + NaOHThe volumes of NaOH in the table are suggested values. Enter your actual volumes of NaOH used in the table (second