PHOTOMETRIC TITRATION OF COPPER (II) WITH EDTA
Experiment
Reading: Harris, Chapter 13 & 6
One result of the commercial availability of modern spectrophotometers is the
development of the photometric titration. In photometric titration, a spectrophotometer is used
so that we can detect an end points graphically. The absorbance of a tritrated species is
measured after each addition of titrant until well past the end point. A plot of absorbance vs.
milliliters of titrant consists of two straight lines that intersect at the end point of the titration.
This method is called the mole ratio method, this and other photometric procedures will be
covered next year in greater depth in Instrumental.
Typical graphs are shown in Figure 1 for the titration of copper (II) with EDTA. The
reaction for this experiment is:
H2Y2- + Cu2+ === CuY2- + 2H+
Where H2Y2- is Na2H2Y
The titration is performed at 625 nm; both the copper-EDTA chelate and the copper (I) ion
absorb at this wavelength, but the molar absorptivity of the chelate is much higher.
Figure 1a is a plot of absorbance (uncorrected for dilution) vs. milliliters of titrant. The
points fall below the extrapolated lines in the end-point region, because the reaction is
incomplete near the equivalence point. After the equivalence point, the added excess EDTA
titrant forces the reaction to completion. The further addition of titrant leads to dilution;
therefore, the absorbance will then decrease slightly.
This dilution effect can be corrected by multiplying the absorbance by a dilution factor,
giving a plot such as that shown in Figure 1b. In practice, a plot such as this can be obtained
without correction by using a large volume (80-10 ml) of solution and a small volume (5 ml or
less) or a relatively concentrated titrant (i.e., dilution effect is small). This will be the procedure
used in this experiment.
The pH is critical for this titration, because a large change in pH changes the effective
binding constant. An acetate buffer is used to maintain the pH between 2.4 and 2.8 to avoid this
problem. This low pH also permits the copper to be titrated in the presence of metal ions that
form weaker complexes with EDTA. Photometric titrations offer additional advantages. For
example, copper (II) ion may be titrated in the presence of an equal concentration of zinc (II) ion,
even thought the copper-EDTA chelate is only slightly more stable (CU-EDTA, log K = 18.8;
Zn-EDTA, log K = 16.5). For the usual indicator or potentiometer titration, the required
difference in log K values is about 4. This difference ensures a favorable end-point equilibrium
in titrating the ion that forms a more stable chelate. The difference is smaller for a photometric
titration for two reasons: (a) the points taken at the beginning of the titration measure only the
copper-EDTA chelate, and (b) the points taken after the end point show no change in absorbance
even though the zinc-EDTA chelate is being formed as the zinc-EDTA chelate is colorless.
EDTA, ml
(a)
EDTA, ml
(b)
Figure 1. Photometric titration of copper (II) with EDTA titrant at pH 2.4: (a) absorbance
uncorrected for dilution; (b) absorbance corrected for dilution.
Procedure
Preparation of Samples
1.
Prepare an acetate buffer solution by adding 4.1 g of anhydrous sodium acetate (or the
equivalent amount of the hydrate) to 50.0 mL of water. Add 1 M hydrochloric acid to the
buffer solution until the pH of the mixture is 2.2, as indicted by a pH meter.
2.
Dry disodium EDTA (Na2H2Y.2H2O) in an oven at 80o C for one hour and then cool in a
desiccator. Prepare 0.2 M EDTA titrant by dissolving about 3.8 grams (known exactly) of
the dry salt and make up to 50.0 mL in a volumetric flask with distilled water. (Note: You
will be preparing a solution that is somewhat difficult to get to dissolve. You might wish
to take advantage of the shake space in the volumetric by not filling completely at the start
of the solvating process.
3.
Quantitatively weigh three sample of copper ore of about 0.9 to 1.1 grams into 250 mL
beakers. To each add 10 mL of 20% nitric acid and cover with a watch glass. Heat the
solution in a hood until the ore dissolves.
4.
Dilute the sample with 25 mL of distilled water and then add 1 M NaOH drop-wise until
the pH reached 3 - 4 on a pH meter.
Data Collection
1.
Turn on the spectrometer to allow it to warm up. Set the wavelength to 625 nm and zero
the meter with distilled water.
2.
Place the beaker on a magnetic stirrer and place a stir bar in the beaker. Add 20 ml of the
acetate buffer solution. At this point, the pH of the solution should be 2.4 to 2.8.
3.
Fill a 50 mL buret with your prepared standard 0.2M EDTA.
4.
Titrate with the 0.2 M EDTA solution in about 0.5 mL increments, allow the solution to
stir for about 1 minute. Take a suitable aliquot from you flask and transfer it to a
spectrometer cell. Record the absorbance. Return your sample aliquot to the beaker and
continue your titration until you are four or five points beyond the endpoint. It is crucial
that you not lose and solution in these transfer processes.
5.
Plot the absorbance of the solution against the volume of the EDTA, determine the
endpoint by extrapolating the two linear portions of the curve to an intersection point. This
can be done using the regression option in Excel. How do you determine the intersection
point. Determine the corresponding volume of EDTA and calculate the percentage of
copper.
6.
Repeat for the last two samples.
What to report
The following should be reported for this lab in the usual format.
1.
2.
Your unknown number! Also collect the following and place into a Table. The number
of mL of EDTA solution required to reach the endpoint for each sample. The weight of
each sample. The percentage of copper in each sample.
Also collect your data into a Table (you will need to do this to find your best fit lines to
determine your end points.)
You will also want to report a best value for the replicate analysis. You will also want to present
your three titration curves as appropriately labeled Figures and show how end points were
evaluated.