Unknown Acid-Base Titration
Introduction
Upon successful completion of this module, the student will be able to describe the
principles of pH measurements and predict the corrosion tendencies of solutions based on
pH. More specifically the student should be able to:
1. Measure the pH of various solutions
2. Estimate the acidity of a solution using colored indicators.
3. Monitor the progress of an acid/base titration with pH sensing electrodes and
graphically determine the end points and equivalence points.
4. Calculate the concentration of an unknown acid.
PART 1: pH and ACIDITY MEASUREMENTS:
The pH, or the concentration of the hydrogen plus ion [H+], is one of the most critical
parameters in predicting if a system will corrode, or determining why a system is
corroding.
pH is defined as the negative log base 10 of the hydrogen plus ion concentration in the
units of moles per liter.
In formula form:
pH =
- log10[ H+ ]
Where :
pH is a number, usually between 0 and 14 for dilute solutions of acids and base in
water.
log10 is the common logarithm to the base 10 (not natural log).
[H+] is the concentration of the hydrogen plus ion in the units of moles per liter.
Please note: The term p in concentration units indicates the -log10 of the concentration of
the indicated ion (i.e. pOH will indicate the -log10[OH+]).
Example Calculation:
If we have 0.038 g of hydrogen plus ion dissolved in 393 mL calculate the concentration
of hydrogen plus ion in moles per liter.
[H+] = (0.038g / (1.0079 g/mol))/(393 mL x 0.001 L/mL)
= 0.096 M = 9.6 x 10-2 M
To calculate the pH of the above solution, substitute into pH = - log 10[H+]
pH
= - log 10[H+]
pH
= -log[9.6 x 10-2]
= 1.02
= 1.0
pH METER
Consider a pH meter to be a black box that is capable of measuring the voltage of a
solution to the nearest millivolt. Although the solution produces a voltage that is directly
related to the concentration of the hydrogen plus ion, measured in moles per liter, it needs
to be electrically connected to the meter via special electrodes. The meter and electrodes
must be calibrated with a solution of accurately known pH, called a buffer and must
compensate for the temperature of the solution. The meter readings are affected by poor
electrical connections as well as static electricity. The following is an example of a
laboratory pH meter.
Figure 1. Typical Laboratory pH Meter
pH ELECTRODES
Two electrodes are required to measure the pH of a solution. One is called the calomel
electrode or voltage reference electrode. The calomel electrode is the reference electrode.
The potential at this electrode remains constant at all hydrogen ion concentrations in the
pH range (i.e. from 1M to 1x10-14). The sensing electrode or the "glass electrode" has a
thin glass dome on its bottom and it is capable of sensing the concentration of the
hydrogen plus ion in solution and generating a corresponding voltage. Typically these
electrodes are housed in one casing but you may see them as separate electrodes. Both
are required for a pH measurement. The potential difference between the two electrodes
is related to pH by the pH Meter usually by using a series of pH standard buffers. Your
pH titration will involve testing of solutions in beakers or flasks. Note that the electrode
must always be supported in a vertical position.
pH INDICATORS
pH indicators are acid/base sensitive organic molecules. They can ionize in solutions of
acids and bases to yield different forms of the parent dye molecules that take on different
colors.
Hin
 H+ + In-
One color in acid form
Another color in base form
At low pH values (high acid (H+) levels) the HIn form of the indicator will predominate
and give a particular color. At high pH values (low acid levels) the In- form of the
indicator will predominate and give a different colors. pH indicators work best in
colorless "water like" solutions. pH indicators do not work well on highly colored,
industrial, sludge type solutions. Many pH indicators exist, and usually the correct one to
determine the pH range you want to measure, can be found.
pH AND TEMPERATURE DEPENDENCE
In theory, at some temperature, the pH of perfectly pure, distilled water should be 7.00,
since the concentration of the H+ ion is 1.00 x 10-7 molar and the solution is considered
neutral.
Consider the following reaction:
H2O  + OHWater breaks apart, or ionizes, to yield one positive hydrogen ion and one negative
hydroxide ion. The concentrations are 1 x 10-7 M and 1 x 10-7 M for H2O, H+ and OHrespectively, at 25oC.
pH OF DISTILLED WATER AT SELECTED TEMPERATURES:
Temperature in oC
Conc. of H+ in [ ]
pH of Solution
0
3.38 x 10-8
7.47
25
1.00 x 10-7
7.00
50
2.24 x 10-7
6.63
100
7.00 x 10-7
6.15
This table clearly shows that pH is temperature dependent. To accurately determine the
pH of water-based solutions, one should first determine the temperature of the solution
and adjust the temperature compensation on the meter correspondingly.
PART 2 - ACID/BASE TITRATION
You now know that an acid has a pH less than 7, a base has a pH greater than 7, and a
neutral solution has a pH equal to 7. Consider a beaker containing a solution of an
unknown dilute concentration of HCl (hydrochloric acid). The pH of the solution would
be less than 7, and could even be as low as 1 or 2. If we added some base, like NaOH, to
the acidic solution, a chemical reaction would occur called neutralization. Some of the
acid and base would react to form a salt, NaCl, (sodium chloride) and H2O (water).
The following reaction shows the results of the neutralization process:
AN ACID AND A BASE REACT TO YIELD A SALT AND WATER
As more and more base was added, there would be a point when the added base would
react with all the acid and there would only be salty water remaining. The pH of the salty
water would be very close to 7. If we were to continue to add more base, the pH of the
resulting solution would end up being much greater than 7, due to the excess base. If we
had an indictor in the solution during the titration, we should have seen one color in the
acid, perhaps another color at neutrality and yet another color at alkaline conditions.
Most indicators indicate only acid and base conditions, but a few often show a third color
at neutrality. If we were to monitor the progress of the titration with pH sensing
electrodes, we would see that in acidic conditions the pH is less than 7, at neutrality the
pH would be equal to 7, and in basic conditions the pH is greater than 7. The rate of
change of pH with respect to volume of base added is slow in acidic conditions, very fast
near neutrality, and very slow in alkaline conditions.
If we were to plot the data, we would get a plot or graph similar to the example in Figure
2.
Figure 2: Example of an Acid Base Titration Curve
CALCULATION OF MOLARITY FOR AN UNKNOWN ACID
Consider the following example:
A technologist pipetted 25.00 mL of an unknown concentration of HCl into a beaker.
The unknown acid was titrated with 37.28 mL of standard sodium hydroxide base with a
concentration 0.1268 moles/liter. The technologist is required to calculate the molarity of
the unknown acid.
Solution:
Neutralization reaction:
The reaction should be balanced first, but in this particular example, it is already balanced
as written. Chemically, the reaction shows that one mole of HCl reacts with one mole of
NaOH to yield one mole of NaCl and one mole of water. At the equivalence point the
original number of moles of unknown acid must equal the moles of known base added.
The equivalence point is given by monitoring the pH verses the volume of standard base
required to neutralize the volume of the acid, of unknown molarity, dispensed.
molarity of acid x volume of acid = molarity of base x volume of base
or
MA x VA = MB x VB
Substituting in our data:
MA
= 0.1891mol/L
Unknown Acid-Base Titration
Procedure
**NOTE: Use instructions provide by the instructor – DO NOT USE INSTRUCTIONS
IN CHEMLAB! To remove the instructions on the screen, and free-up more working
area, perform the following operation: click on the OPTIONS tab; then click on LAB
ONLY. The instructions “disappear” and all of the area is now lab space.**
1.
pH MEASUREMENT BY METER
Step 1: Obtain a 100mL beaker and equip with pH meter (right click on beaker → pH
Meter)
Step 2: Obtain sample unknown ‘A’ to the pH of, and place 50 mL in the beaker.
Step 3: While gently stirring, determine and record the pH of the solution.
Step 4: Repeat Steps 1 – 3 for unknown ‘B’ and unknown ‘C’.
2.
ACIDITY MEASUREMENT WITH INDICATORS
Step 1: Set up two 50 mL test tubes.
Step 2: Label one as an Acid “A” and one as a Base “B”
Step 3: Place 35 mL of 0.100M HCL into the Acid test tube and 0.100M NaOH in the
Base test tube.
Step 4: Add 5 drops of an phenolphthalein into each test tube.
Step 5: Record the resultant color for the acid and base.
Step 6: Repeat steps 1 – 5 for the other 3 Acid - Base Indicators
3.
pH MONITORED ACID/BASE TITRATION
Step 1: Obtain acid, in a 100 ml Erlenmeyer flask add 35 ml of an Unknown HCl
solution. (Choose either unknown ‘A’, ‘B’, or ‘C’)
Step 2: Add an indicator to the acid, select the flask and add 2 drops of phenolphthalein
indicator. The indicator menu is available under the chemicals main menu
(ChemicalsIndicators) or the context menu.
Step 3: Fill buret with NaOH, obtain a 50 ml buret and fill with .100M NaOH solution.
Step 4: Place Unknownl solution on stir plate, Select Equipment→ Hot plate and mag stir
– place the Erlenmeyer flask on the plate and turn on the magnetic stir bar.
Step 5: Use ChemLab to collect data and generate titration curve



Add a pH meter to the acid solution.
Turn on the collection of titration data for the Erlenmeyer flask by selecting the
"collect titration data" menu from the procedures menu (or from the right-mouse
context menu).
Open the Titration data window by selecting the "View titration data" from the
Procedures menu.
Step 6: Titrate NaOH into HCl until end point, record initial buret volume and add NaOH
(quickly at first then slowly) until the HCl solution turns pink and record the final buret
volume of NaOH in buret.
Step 7: Copy the LABELED titration curve from the Titration data window to Word,
Use the copy from the Titration data window's edit menu and paste into Word – print
your titration curve. Indicate the equivalence point on your graph. Indicate the
equivalence point volume on your graph.
Unknown Acid-Base Titration Name:__________
Section:_________
Observations
Data
1.
pH MEASUREMENT BY METER
Unknown A
Unknown B
Unknown C
2.
pH = _______
pH = _______
pH = _______
ACIDITY MEASUREMENT WITH INDICATORS
Indicator
Color in Acid Solution
Phenolphthalein
Methyl Red
Phenolsulfonephthalein
(Phenol Red)
Bromothlmol Blue
3.
Color in Basic Solution
_____________
_____________
_____________
_____________
_____________
_____________
_____________
_____________
pH MONITORED ACID/BASE TITRATION
UNKNOWN ACID LABEL INFORMATION:
(circle one) ‘A’ – ‘B’ – ‘C’
CONCENTRATION OF STANDARDIZED NaOH:
__0.100 mol/L
Equivalence point volume from graph:
_________mL
Calculations
Molarity of unknown acid :
(use the example provided in the Introduction for guidance)
________[M]
Questions
acid/base titration questions
1.
Why should the solution in the beaker be stirred continuously throughout the
titration?
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
__________________________________________________________________
pH & Acidity Measurement Questions
1.
Calculate the molarity of the resulting solution if 0.126 grams of Hydrogen ion
[H+] is dissolved in 250 mL of distilled water.
________[M]
2.
Calculate the pH of the solution considered in question 1.
(use the example provided in the Introduction for guidance)
pH = _______
3.
Concentrated hydrochloric acid (HCl), or muriatic acid as it is called in industry,
is available in a concentration of 12 M. It ionizes to yield one acid proton per
molecule of HCl. Calculate the pH of this very strong and concentrated acid.
pH = _______
Discussion