EXPERIMENT 5
ACID-BASE TITRATION
INTRODUCTION
Much of chemistry and biology is concerned with the behavior of acids and bases. Acids and bases
are participants in many reactions in nature, and many reactions require a particular level of acidity or
basicity. For example, the concentration of acid in the stomach must be quite high, in comparison to the
rest of the body, to break down proteins during digestion.
This experiment involves the determination of the amount of acid or base in an aqueous solution,
using an acid-base titration. A titration is the addition of a carefully measured volume of a solution of one
substance to a solution of another substance until a reaction between the two substances is just complete.
The titration in this experiment is based on a fundamental chemical property of acids and bases: their
neutralization reaction producing water and a salt:
HA + BOH
H2O + BA
(5-1)
2 H2O(l)
(5-2)
for which the net ionic equation is
H3O+(aq) + OH-(aq)
If the amount of acid in one solution is known, then the amount of base in a second solution can be
determined by measuring how much of it is required to neutralize the acid. Similarly, if the amount of
base in one solution is known, the amount of acid in another solution can be determined by a titration.
EQUATION 5-1 can be used as a general equation to represent acid-base neutralization reactions.
However, the generalized formulas in it do not accurately represent all acids and bases, for some acids
contain more than one reactive hydrogen atom per formula unit and some bases contain (or produce) more
than one hydroxide ion per formula unit; examples are H2SO4, H3PO4, Mg(OH)2, and Fe(OH)3.
EQUATION 5-2 is a more general equation; it shows that a neutralization reaction involves equimolar
amounts of hydrogen ions (in aqueous solution as hydronium ions) from the acid and hydroxide ions from
the base:
The term equivalent mass emphasizes the essential chemical behavior of acids and bases. One
equivalent mass of an acid is the mass of the acid that contains one mole of ionizable hydrogen atoms.
One equivalent mass of a base is the mass of the base that produces one mole of hydroxide ions. For some
compounds, equivalent mass equals molar mass; examples are HCl, HC 2H3O2, NaOH, and NH3. For other
compounds, equivalent mass is a particular fraction of the molar mass. For example, one mole of
Ca(OH)2 contains two moles of hydroxide ions and so one equivalent = 1/2 mole = 37 g. For H 3PO4, one
equivalent = 1/3 mole = 33g. In each case:
equivalent mass =
molar mass
number of H or OH per formula unit
+
EXPERIMENT 5
(5-3)
5-1
In an acid-base titration, the equivalence point is reached when equimolar amounts of hydrogen ion
from acid and hydroxide ion from base have been combined in solution, EQUATION 5-2. We can also say
that at the equivalence point, equal numbers of equivalent masses (or equivalents) of acid and base have
been combined. If the identity of one reactant is known, the number of equivalents of that reactant can be
calculated. The titration requires the same number of equivalents of the other reactant, and so the number
of equivalents of it is known also. Since you will be titrating an unknown acid with a solution of sodium
hydroxide,
at the equivalence point, equivalents of H3O+ = moles of OH−
(5-4)
TECHNIQUE
You will carry out acid-base titrations using the acid-base indicator phenolphthalein. It is colorless in
acidic solutions and pink in basic solutions; the color ranges from pale pink to dark purplish pink,
depending on the concentrations of indicator and base. If you titrate to the appearance of a light pink
color, the amount of excess base will be numerically insignificant; hence, this lightest possible pink color
indicates the end point. By choosing an indicator properly, endpoint = equivalence point.
Your first activity will be to standardize a solution of NaOH; this means that you determine its
concentration as accurately and precisely as possible. The NaOH solution supplied is labeled as having a
concentration of 0.1M; you will determine the actual concentration to at least three significant figures.
You do this by using the NaOH to neutralize an accurately weighed sample of KHC8H4O4 (potassium
hydrogen phthalate, KHP), a monoprotic acid that has a molar mass of 204.22 g. You can calculate the
number of moles of acid used in the titration:
Moles KHP = g KHP/molar mass KHP
(5-5)
Since KHP is monoprotic, this is the same as the number of moles of hydrogen ion provided by the
acid. It is also the number of moles of hydroxide ions from the base that the acid neutralizes. The NaOH
solution is dispensed from a buret and its volume is carefully measured; you use the numbers of moles of
hydroxide ions and volume of NaOH solution to calculate a precise molarity (moles/liter) of the NaOH
solution.
You will use your standardized NaOH solution to determine the equivalent mass of an unknown solid
acid. Since you do not know whether your unknown acid is mono- di- or triprotic, you will generate
enough data to determine the acid’s equivalent mass, but not its molar mass. You do this by using your
standardized NaOH solution to neutralize a carefully weighed sample of the acid. Using the molarity and
volume of the NaOH solution, you can calculate moles of hydroxide ions required to neutralize the acid.
At the equivalence point of the titration, this equals equivalents of hydrogen ions provided by the acid,
EQUATION 5-4. The equivalent mass of the acid is the mass of the acid that contains one mole of
ionizable hydrogen. Therefore, you can calculate the equivalent mass from your data:
equivalent mass unknown acid = g unknown acid/equivalents of H3O+
EXPERIMENT 5
(5-6)
5-2
EQUIPMENT NEEDED
analytical balance
beakers
buret
buret clamp
Erlenmeyer flasks
graduated cylinders
CHEMICALS NEEDED
KHC8H4O4; potassium hydrogen phthalate (KHP)
0.1M NaOH; sodium hydroxide
phenolphthalein indicator solution
unknown solid acid
PROCEDURE
A. Standardization of NaOH Solution
1. Obtain ~100 mL of 0.1 M NaOH solution in a beaker. Clean a buret, rinse the buret with a small
amount of the solution, and then fill it with the NaOH solution. Be sure to drain some of the solution
so that the solution completely fills the tip of the buret.
2. On the analytical balance, weigh out 0.2–0.3 g of KHC8H4O4 (be sure to record to the nearest 0.0001
g) and put it into an Erlenmeyer flask. To do this, first place a small beaker or weighing boat on the
analytical balance, and press the bar or button on the balance to zero the balance. Then add 0.2–0.3 g
of the KHP to this container. Go back to your laboratory desk and pour the acid into the flask; use a
stream of water from a wash bottle to complete the transfer. Dissolve the acid in about 20 mL of
distilled water. Add two drops of phenolphthalein indicator solution to this.
3. Record the initial buret reading to the nearest 0.01 mL. If there is a drop of solution on the tip of the
buret, remove it by touching the tip of the buret to the inside wall of a waste beaker. Place the flask
(containing the acid) under the buret and place a piece of white paper under the flask so that you can
easily see any color in the solution. Position the buret so that the tip of it is just slightly below the rim
of the flask.
4. Begin the titration by adding several mL of base solution in a rapid stream; constantly swirl the flask.
At the site of addition, there will be a pink color that disappears quickly as the acid and base solutions
mix. When this color seems to linger, turn the stopcock to shut off the flow of solution from the buret.
At this point, you can introduce the NaOH solution into the flask by turning the stopcock 180
degrees. The speed with which you do this will determine how much solution is allowed into the
flask—a leisurely turn will introduce a few tenths of a milliliter, while a quick flick will introduce a
single drop. Decrease the rate of addition of base as the pink color persists longer before it disappears
(remember to swirl after each addition of solution). Occasionally wash solution off the inside wall of
the flask with a stream of water from a wash bottle. Add base until the lightest possible pink color
persists throughout the solution for at least 20 seconds. If the solution has a dark pink color, you
EXPERIMENT 5
5-3
probably passed the end point; be sure to record this information to help you decide later about the
relative accuracy of the titrations. Record the final buret reading (to the nearest 0.01 mL) about a
minute after completing the titration; this time lag allows solution on the inside wall of the buret to
flow into the bulk of the solution.
5. Before you begin the next titration of a weighed KHC8H4O4 sample, use base volume and acid mass
from the first titration as a conversion factor to approximate the volume of base that this mass of acid
will require. Add base rapidly until you approach this volume and then add base slowly. Keep in mind
that if you over-titrated the first sample, this approximate volume will probably be high.
6. Carry out at least three titrations. If you know that you over-titrated a particular acid sample, or made
some other mistake that would affect the accuracy of your results, you should perform another
titration. You will include your best three trials in your final report. It is recommended that you
calculate the molarity of the NaOH solution for each trial while you are in the lab—that way you can
assess the reproducibility of your titrations (the molarities should agree within ~5%)
B. Titration of Unknown Acid
7. Obtain an unknown solid acid and record its identification code. On the analytical balance, weigh out
0.1–0.2 g of the acid (be sure to record to the nearest 0.0001 g) and put it into an Erlenmeyer flask.
Dissolve the acid in about 20 mL of distilled water and add two drops of phenolphthalein solution.
Titrate the acid to the pale pink end point with your standardized NaOH solution.
8. If this first titration requires less than 10.00 mL of base solution, you will have only three significant
figures in your calculations. For the next titrations, increase the acid sample size so that the titration
will require at least 10.00 mL of base, but less than the capacity of the buret. If the first titration
requires more base than the capacity of the buret, refill the buret, record the buret reading, and
continue the titration. For the next titrations, decrease the acid sample size so that the titration will
require less base solution than the capacity of the buret, but at least 10.00 mL.
9. Carry out at least three titrations that require a volume of base that is at least 10.00 mL but less than
the capacity of the buret. If you suspect that one of your titrations is inaccurate, and if time permits,
perform an additional titration. You will include your best three trials in your final report.
10. All the solutions from your titrations may be disposed of by combining the solutions and washing
them down the drain. Dispose of leftover solid unknown acid in the trash can.
EXPERIMENT 5
5-4
EXPERIMENT 5
REPORT SHEET
Name: _______________________________________
Date:__________
A. STANDARDIZATION OF NaOH SOLUTION
TRIAL 1
TRIAL 2
TRIAL 3
TRIAL 4
mass of KHC8H4O4
__________
__________
__________
__________
moles of KHC8H4O4
__________
__________
__________
__________
__________
__________
__________
__________
initial buret reading
__________
__________
__________
__________
final buret reading
__________
__________
__________
__________
vol. of NaOH
__________
__________
__________
__________
molarity of NaOH
__________
__________
__________
__________
sample calculation:
moles of NaOH
sample calculation:
sample calculation:
avg. M of NaOH
_______________________
EXPERIMENT 5
5-5
B. TITRATION OF UNKNOWN ACID
Unknown identification code ___________
Average Molarity of NaOH from Part A. _________
TRIAL 1
TRIAL 2
TRIAL 3
TRIAL 4
mass of acid
__________
__________
__________
__________
initial buret reading
__________
__________
__________
__________
final buret reading
__________
__________
__________
__________
vol. of NaOH
__________
__________
__________
__________
moles of OH
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
sample calculation:
equivalents of H3O+
sample calculation:
equivalent mass of acid
sample calculation:
avg. equiv. mass of acid
_______________________
EXPERIMENT 5
5-6
Notes
Experiment 5
You will work individually on this experiment.
Obtain about 100 mL of the 0.1M NaOH solution.
You should share a buret clamp with a neighbor.
Potassium hydrogen phthalate is commonly abbreviated as KHP. This may be what
you will see on the label of the reagent jar.
Remember to read and record your volume measurement to the nearest 0.01 mL!
Don’t forget to write down the unknown number for your unknown acid!
Next Week: You will be performing an experiment next lab period that will involve
uploading your data to the Chem21 website. In order to do this, you have
been assigned a Student ID number which can be found on the main page
(right after logging in) at the Chem21 website. You must bring this
Student ID number with you to lab the day you are to perform Exp.
12.
10/10
EXPERIMENT 5
5-7