Acid-Base Titration Lab
Introduction:
You are a chemistry student at the Whoopsacola Institute of Technology (WIT). As
part of your General Chemistry course, you have agreed to set up a lab utilizing sodium
hydroxide solution. Your professor tells you that the solution has already been made, but he
doesn’t remember the concentration of the base, but it is vital to the success of the lab. He
says not to worry about it, that he put a label on the flask containing the sodium hydroxide
solution with the formula, concentration and preparation date on it. Satisfied, you leave him
to prepare the lab. You first take an extended coffee break and return to the lab, only to
find that some prankster has crossed out the concentration on the label, the professor is
nowhere to be found and the lab is to start in half an hour. After regaining consciousness,
you review your options. You may 1) drop out of college while you can, or 2) perform a
titration of the NaOH using the bottle of standardized HCl conveniently nearby. Being a
conscientious chemistry student, you opt for plan 2.
Acids and bases neutralize each other. When both are mixed so that their hydrogen
ion concentration and hydroxide ion concentration are equal, the pH will equal 7, and the
solution will be neutral. This is called the equivalence point, because the concentrations of H+
and OH- are equal to each other. The concentration of base may be determined by titration,
which is a process that involves adding the base to a quantity of acid of known concentration
that has an indicator added to it. In this case, phenolphthalein is the indicator of choice, for
it is colorless in acidic solutions, but brilliant magenta in alkaline (basic) solutions. At first,
the solution of HCl is entirely acidic, but as NaOH is added, the alkalinity increases. When
the concentration of OH- ions equals that of the H+ ions, you have reached neutralization,
also called the equivalence point. At roughly this point, if any more NaOH is added, the
presence of additional OH- ions will cause the phenolphthalein indicator to turn pink
throughout the solution. This is called the endpoint, the point at which the indicator turns
color. At this point, the solution will be slightly alkaline (basic), so the paler the pink color is,
the closer the endpoint is to the equivalence point, and the more accurate your results.
The number of moles of H+ equals the number of moles of OH- when the solution is
neutral. At this point, the titration is complete. Therefore, we can write moles H + = moles
OH-. Since we are dealing with concentrations in molarity (moles/L), moles = M X V (molarity
X volume). Therefore, the equation is derived:
(#H+)MaVa = (#OH-)MbVb
Variable
#H
Ma
Va
+
Acid
What it means
+
Number of hydrogen (H ) ions in
the formula of the acid
Molarity of the acid
Volume of the acid
Base
What it means
Variable
#OH1
Mb
Vb
Number of hydroxide (OH1) ions in
the formula of the base
Molarity of the base
Volume of the base
© 2005, Mark Rosengarten, Pam Iacovella
1
So, to solve for the molarity of the base, rearrange the formula to solve for Mb, plug in the
numbers and voila! Molarity of base is found! You will do four trials, calculate the molarity of
each trial and then calculate the average molarity.
Materials:
Ringstand
buret clamp
acid buret
base buret
2 funnels
white paper
2 beakers (one for acid and one for the
250 mL Erlenmeyer flask
phenolphthalein
base)
0.100 M standard HCl solution
NaOH solution of unknown concentration
Procedure:
1) Read and record in the data table, the initial volume of both the acid and base burets.
2) Let 10.0 mL of the acid flow into a 250 mL Erlenmeyer flask. Add to it three drops of
phenolphthalein.
3) Add to the flask NaOH from the other buret, placing a piece of white paper under the
buret to make the pink color easier to see. Swirl the flask as you add the base. The pink
color will disappear. The more base that is added, the longer it will take for the pink color
to disappear. Decrease the rate of adding the base until you are adding it one drop at a
time. When one drop is enough to turn the solution pink for 15 seconds, the solution has
reached the endpoint. If you accidentally go past the endpoint, add additional acid to the
flask until it clears up, then add the base slower until a pale pink color persists for 15
seconds.
4) If too much base is added (giving a deep pink color), then go back and clear up the solution
with acid one drop at a time, then try the base again.
5) Record the final volume of acid and base indicated by the burets.
6) Pour the solution in the 250-mL Erlenmeyer flask into the sink and rinse the flask out
several times. Dry the flask.
7) Repeat steps 2-6 three more times. The initial volume of the next trial will be the final
volume of the previous trial. This is done to conserve acid and base.
Example:
You will have enough of each solution in the burets to accomplish three titrations. Discard
the neutralized solutions before going on to the next titration. Since you know how many mL
of base were needed to neutralize the acid in the first trial, you can let the base run out fast
until it gets to about 2 mL before the endpoint. Record the final volume indicated on the
burets after each trial.
SEE TEACHER FOR CLEAN–UP INSTRUCTIONS
© 2005, Mark Rosengarten, Pam Iacovella
2
Name:
Lab Partners:
Period
Date Due:
Acid-Base Titration
Data:
Trial 1
HCl
NaOH
Trial 2
HCl
NaOH
Trial 3
HCl
NaOH
Trial 4
HCl
NaOH
Initial Volume (mL)
Final Volume (mL)
Volume Used (mL)
(use in calculation)
Analysis:
1) Calculate the volume of acid used and the volume of base used for each of the four trials
and place it in the bottom row of the data table.
2) Given the molarity of the acid (see materials list) and your values for volume used of acid
and base, calculate the molarity of the base for each trial. Show all work, including the
formula setup and units. Put your final answer in the right-hand box for each trial.
Trial 1
Show Work
Molarity
of Base
Show Work
Molarity
of Base
Trial 2
© 2005, Mark Rosengarten, Pam Iacovella
3
Trial 3
Show Work
Molarity
of Base
Show Work
Molarity
of Base
Show Work
Average
Molarity
of Base
Trial 4
Average of the 4 trials
Questions
1) If you were to round the calculated molarities for all four trials and the average molarity
to one significant figure, how similar are your results? (Complete sentences!)
_____________________________________________________________________
__________________________________________________________________
___________________________________________________________________
© 2005, Mark Rosengarten, Pam Iacovella
4
2) The accepted value for the molarity of the base is 0.200 M. Calculate your percent error,
showing all work. Use your average molarity as your experimental value. (Math formula, all
numbers with units and answer rounded to proper significant figures required!)
3) If too much base was added to the flask during the titration, this problem could be
remedied by adding more acid to the flask until the desired shade of pink was obtained.
Why would this remedy not affect the outcome (the calculation of the molarity of the
base) in this experiment? Remember what was being made equal in this lab. (Complete
sentences!)
_____________________________________________________________________
__________________________________________________________________
_____________________________________________________________________
_____________________________________________________________________
4) If 30.0 mL of a 0.500 M KOH solution are needed to neutralize 10.0 mL of HCl of unknown
concentration, what is the molarity of the HCl? (Math formula, all numbers with units and
answer rounded to proper significant figures required!)
5) How many mL of 0.100 M NaOH are needed to titrate 20.0 mL of 0.100 M H2SO4?
(Math formula, all numbers with units and answer rounded to proper significant figures
required!)
© 2005, Mark Rosengarten, Pam Iacovella
5