Generating a Titration Curve
Strong Acid & Strong Base

In this exercise you are going to generate a plot of the changes in pH as a strong acid is added to a strong base.  This curve is called a titration curve.

Titration is the progressive addition of measured volumes of a solution of known concentration to an unknown substance.  The reaction reaches an observable endpoint, where the amount of  unknown is calculated from the volume and concentration.  In an acid-base titration the chemical relationship is that 1 mole of H3O+ can neutralize 1 mole of OH-, and acid-base indicator dyes are often used to detect the endpoint.

In this exercise you are going to construct a spreadsheet for the titration of a strong acid by a strong base.  Here are the assumptions to use in constructing this spreadsheet:

Then

  • the moles of  H3O+ originally present = [Acid] x Lacid
  • the moles of added OH- = [Base] x  Lbase

 

These statements are true because the formula for molar concentration is  [Molar Concentration] = mol/L so mol = [Molar Concentration] x L and these are strong acids and bases

 

This is true until the point when exactly equal moles of acid and base have been added.  At this point, the pH = 7, since it is a neutral solution.  From this point on, if more base is added, it is equivalent to adding base to a solution of pure water.  Therefore:

Procedure:

1. Create a new Excel spreadsheet.  Add a section to it to identify the spreadsheet, and to include the initial concentrations like this: 
 

Click here to load a prepared template

 
Now label the cells in the spreadsheet as follows:

2. In column A, enter the amount of added base, starting with 0 ml, and increasing by 5 mL increments to 100 mL.
3. In column B, enter a formula that will calculate pH, ( -log[H3O+]) from the negative log of column I.
4. In column C, calculate the volume of base in Litres, by creating a formula that is related to the value in column A.
5. In column D, calculate the volume of acid used, by creating a formula that is related to the value in cell C4.  Note: This formula must not change when copied into each cell of column D.
6. In column E, sum the values in column C and D.
7. In column F, calculate the mol of OH-, using the value in cell C6 (which must not change when copied into each cell) and the volume (in Litres) in column C.
8. In column G, calculate the original mol of H3O+ using the value in cell C5 (which must not change when copied into each cell) and the volume (in Litres) in column D.
9. In column H, calculate the original mol of H3O+ using the value in cell C5 (which must not change when copied into each cell) and the volume (in Litres) in column D.
10. In column I, calculate the [H3O+] by dividing the mol H3O+ (column H) by the total volume (column E).
11. In column J, calculate the [OH-] from the Kw formula.
12. Copy these formulas down each column, until you reach the row where the mL of Base = mL of Acid.  At this point the above formulas no longer work, because the endpoint has been reached.  In the resulting 100 mL of neutral water, there will be 1 x 10-8 mol of H3O+ so enter that value manually into cell H20.

13. Copy the existing formulas for columns B through G for the entire spreadsheet.
14. From the point of neutrality on (row 21 and on) leave column H blank.
15. In cell J21 enter a formula that will calculate the excess [OH-], by subtracting (column F - column G) and dividing by the total volume (column E).  Copy this formula into each following cell in the column.
16. In cell I21, enter a formula that will calculate the [H3O+] from the Kw formula.
17. Highlight the data in columns A and B and use the chart wizard to plot a graph of pH as as function of added volume of base.
18. Notice on the graph how steeply the pH changes near the middle.  In fact it changes even more steeply than is shown, because of the large volume increment we have used.  Fix your spreadsheet by:

When satisfied with your spreadsheet, print the graph.
19. Compare your graph to the acid-base indicator dyes and determine which indicators would be acceptable, based on the idea that the indicator must change color near the equivalence point for the reaction.  Because of the rapid change in pH near the equivalence point, several indicators should give almost identical results.



Copyright © 1998 - 2008 David Dice