# Example Problem: Determining pH in a Titration## ObjectiveDetermine the pH at the point in the titration of 40.0 mL of 0.200 M hydrazine solution (H₂NNH₂) with 0.100 M nitric acid (HNO₃) after 80.0 mL of the strong acid has been added. The Kb value for H₂NNH₂ is 3.0 × 10⁻⁶.## ProcedureBased on the result of the acid-base reaction, set up the ICE (Initial, Change, Equilibrium) table in order to determine the unknown concentrations.### Reaction Equation\[ \text{H}_2\text{NNH}_3^+(aq) + \text{H}_2\text{O}(l) \rightleftharpoons \text{H}_3\text{O}^+(aq) + \text{H}_2\text{NNH}_2(aq) \]## ICE Table Format| \[ \] | \[\text{H}_2\text{NNH}_3^+(aq)\] | \[\text{H}_2\text{O}(l)\] | \[\text{H}_3\text{O}^+(aq)\] | \[\text{H}_2\text{NNH}_2(aq)\] ||-------|-----------------|---------------|-----------------|---------------|| **Initial (M)** | \[ \] | \[ \] | \[ \] | \[ \] || **Change (M)** | \[ \] | \[ \] | \[ \] | \[ \] || **Equilibrium (M)** | \[ \] | \[ \] | \[ \] | \[ \] |## Graphical ControlsAt the bottom, select from the pre-encoded concentration fields:- Initial Concentrations: - 0.100 - 0.200 - 0.0333 - 0.0667- Changes in Concentration: - \( +x \) - \( -x \)Final Equilibrium Concentrations:- \[ 0.100 + x \]- \[ 0.100 - x \]- \[ 0.200 + **Titration Question: Determining pH****Problem Statement:**Determine the pH at the point in the titration of 40.0 mL of 0.200 M H₂NNH₂ with 0.100 M HNO₃ after 80.0 mL of the strong acid has been added. The value of Kb for H₂NNH₂ is 3.0 x 10⁻⁶.**Instructions:**Use the table below to determine the moles of reactant and product after the reaction of the acid and base.**Reaction Table:**| | H₂NNH₂(aq) | + | H⁺(aq) | → | H₂NNH₃⁺(aq) ||-------------------------------|-----------|---|--------|---|-------------|| **Before (mol)** | | | | | || **Change (mol)** | | | | | || **After (mol)** | | | | | |**Interactive Options:**- Below the table, there are several options with different values and expressions to select from, such as: - 0 - 0.100 - +x - -x - Numerical values like 1.00 x 10⁻³, 2.00 x 10⁻³, etc. **Note:**Make sure to use the given values alongside the expressions and equations to solve for the unknowns in the reaction table. Adjustments can be made using the "RESET" button if necessary.

given values : Kb =3×10-6pKb=-log (3×10-6) =5.52 for weak base hydrazine : molarity = 0.2 M…

Determine the pH at the point in the titration of 40.0 mL of 0.200 M H2NNH2 with 0.100 M HNO3 after 80.0 mL of the strong acid has been added. The value of Kb for H2NNH2 is 3.0 x 106. 1 3 4 NEXT Use the table below to determine the moles of reactant and product ater the reaction the acid and base. H2NNH2(aq) H*(aq) H,NNH; *(aq) + Before (mol) Change (mol) After (mol)

Determine the pH at the point in the titration of 40.0 mL of 0.200 M H2NNH2 with 0.100 M HNO3 after 80.0 mL of the strong acid has been added. The value of Kb for H2NNH2 is 3.0 x 106. 1 3 4 NEXT Use the table below to determine the moles of reactant and product ater the reaction the acid and base. H2NNH2(aq) H(aq) H,NNH; (aq) + Before (mol) Change (mol) After (mol)

Chemistry

10th Edition

ISBN:9781305957404

Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Chapter1: Chemical Foundations

Section: Chapter Questions

Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...

See similar textbooks

Similar questions

30.00 mL sample of a 0.200 M (aq) solution of a weak base is titrated with a 0.250 M HCl at 25 degrees celcius. Find the pH at each volume of titrant added. The kb for the base is 2.6*10^-6 1. 10.00 mL 2. 15.00 mL 3. 30.00 mL
Consider the titration of 40.0 mL of 0.200 mol/L HCOONa (aq) with 0.300 mol/L HCI (aq). Ka for HCOOH is 1.8×104. Determine the pH of the original solution. Use the given values to set up an ICE table. Determine the unknown equilibrium concentrations. Initial (M) Change (M) Equilibrium (M) 1 HCOONa(aq) H₂O(1) HCOOH(aq) OH (aq) 5 RESET
An analytical chemist is titrating 143.6 mL of a 0.7000M solution of piperidine (C,H10NH) with a 0.4000M solution of HNO3. The p K, of piperidine is 2.89. Calculate the pH of the base solution after the chemist has added 267.7 mL of the HNO, solution to it. Note for advanced students: you may assume the final volume equals the initial volume of the solution plus the volume of HNO, solution added. Round your answer to 2 decimal places. pH
When titrating 125 mL of nitric acid HNO3 0.165 mol/L, with KOH 0.219 mol/L.a) What is the pH of the solution before starting the titration, and what volume of KOH (in mL) is required to reach the equivalence point?b) What is the pH of the solution after the addition of 44.5 mL of KOH solution?c) What is the pH of the solution at the equivalence point, and the pH of the solution after the addition of 200 mL of KOH solution?
NaOH: 0.100 M Molarities of Solutions HCI: 0.1007 C2H3O2: 0.1005 M M-HC2H3O2: Part 1 - Titration of Hydrochloric Acid becubong ovisnotsuite griau bloe piles to solar mob of Using your titration curve, what is the mL base needed to reach the equivalence point? CORDAS Jag tog nolluloa bien ne bs agnied bio 25 nulov aumey He To Joig leciiaste axem uov umLoitsu (A clients es wolle ndog conalovispa en bron Han qe For each point shown below in the titration, record your measured pH from your data, then calculate the pH of the solution. Show your work for all calculations in the space provided. fomento orfT [HC1] = [H+] anos evig illw blos Acow a bris (10moq nobsilit 1. What was the pH of the hydrochloric acid before any base was added? -log[H+] = -log[0.1007] Measured pH = 4.1 pH = -log [H+] where [HCI] = [H+] Calculated pH: = 1.0 06 Hq =0.99697 ~1.0 2. What was the pH after 10.0 mL of sodium hydroxide solution was added? Measured pH = 1.3 Volume of HCI solution in liters: 002 0.01 net Um L…
An analytical chemist is titrating 64.8mL of a 0.3000M solution of ammonia NH3 with a 0.7300M solution of HNO3 . The pKb of ammonia is 4.74 . Calculate the pH of the base solution after the chemist has added 31.6mL of the HNO3 solution to it. Note for advanced students: you may assume the final volume equals the initial volume of the solution plus the volume of HNO3 solution added. Round your answer to 2 decimal places.
In a titration, if 12.00 mL of vinegar requires 19.85 mL of 0.4500 M NaOH to reach the equivalence point. What is the molarity of the vinegar solution? Hint: Vinegar contains acetic acid (HC2H3O2).
a buffer solution is often encountered during the titration of aweak acid. In such a titration, there is a strong base (often sodium hydroxide, as in today’s lab)which is being added to the weak acid. When the strong base reacts with the weak acid, theresult is the conjugate base of the weak acid. It is essential that you not confuse these twobases during the discussion below, and that you write your report so that it is clear which baseyou are talking about. If the pH of the acid solution is monitored during the titration, a pHprofile like the one below can be plotted. For monoprotic acids it will be sigmoid in shape:The Henderson-Hasselbalch equation helps to make sense of this curve (the base referredto is the conjugate base of the weak acid).pH = pKa + log ([base]/[acid])If calculations are desired, two points are particularly important. The first, at the steepest pointof the graph, is the equivalence point. At that point the acid has been completely consumed bythe strong base…
A 76.0 mL sample of 0.0100 M HIO4 is titrated with 0.0200 M RbOH solution. Calculate the pH after the following volumes of base have been added. (a) 14.1 mL pH = (d) 39.9 mL pH = (b) 37.2 mL pH = (e) 69.5 mL pH = (c) 38.0 mL pH =
Help please
propionic acid, ?a=1.34×10−5,Ka=1.34×10−5, 3.00 M sodium propionate, CH3CH2COONa 1. The final volume of buffer solution must be 100.00 mL and the final concentration of the weak acid must be 0.100 M. Based on this information, what mass of solid conjugate base should the student weigh out to make the buffer solution with a pH of 5.00? ( in g) 2. Based on this information, what volume of acid should the student measure to make the buffer solution? (in ml)
A sample of acetic acid is titrated with a standardized NaOH solution. Before the end-point of the titration, which of the following must be true?