TITRATION
•
This involves removing small samples
from the reaction mixture at different times
and then titrating the sample to determine
the concentration of either one of the
reactants or one of the products at this
time. Then we can draw concentration vs
time graph.
•
This method is suitable for quite slow reactions
so that the time taken to titrate the mixture is
insignificant compared to the total time taken
for the reaction.
• We can overcome this difficulty by quenching
(stopping) the reaction before carrying out the
titration.
• This can be done:
1. Cooling the reaction mixture rapidly to a very
low temperature
2. Adding an excess amount of a compound that
reacts rapidly with one of the reactants.
• This method can be suitable for the reaction:
H2O2(aq) + 2H+(aq) + 2I-(aq)
2H2O(l) + I2(aq)
The amount of iodine produced can be measured
by titrating the mixture with aqueous sodium
thiosulfate. The reaction can be quenched by
adding excess insoluble solid base to neutralize
the acid that is required for the reaction.
In analytical chemistry, the most important use comes from
the fact that the thiosulfate anion reacts stoichiometrically
with iodine, reducing it to iodide as it is oxidized to
tetrathionate:
2 S2O32−(aq) + I2(aq) → S4O62−(aq) + 2 I−(aq)
TITRATION CALCULATIONS
• Titration is a technique which involves
measuring the volume of one solution
which just reacts completely with another
solution.
• One of the solutions have an accurately
known concentration (standard solution)
and this is used to find the concentration
of the other solution.
• The volume of one of the solutions will be
measured with a pipette and put into a conical
flask and a few drops of an indicator is added.
pipette
Conical flask
Second solution is run in from a burette until the
indicator just changes color.
Indicator: a
chemical
substance that
changes colour
at the
equivalence
point.
Burette
Indicator
Burette has a tap
and is calibrated so
as to measure a
variable volume of a
liquid.
The equivalence point
is the point where the
moles of titrant added
equals the moles of
substance initially in the
solution being titrated.
Using a Burette
When reading a buret it is important that
your line of sight be in a direction
perpendicular to the buret column. You
will read the volume from the bottom of the
meniscus.
The use of a buret card prevents the
lightenning of the bottom of the meniscus by
random reflections in the laboratory.
A 50 mL buret can be read to
±0.01 mL
Doing a Titration
Begin by preparing your buret. Your
buret should be filled with titrant solution. You
should check for air bubbles and leaks,
before proceding with the titration.
Take an initial volume reading and
record it in your notebook.
Prepare the solution to be
analyzed by placing it in a clean
erlenmeyer flask and add the
indicator.
Use the buret to deliver titrant into
erlenmayer flask. You will see the indicator
change color when the titrant hits the solution
in the flask, but the color change disappears
upon stirring.
Approach the endpoint more
slowly and watch the color of your flask
carefully and be sure all titrant is mixed
in the flask.
Make sure you know what the endpoint
should look like. For phenolphthalein, the endpoint is
the first permanent pale pink.
If you think you might have reached the
endpoint, you can record the volume reading and add
another partial drop.
When you have reached the
endpoint, read the final volume in the buret
and record it in your notebook.
Subtract the initial volume to
determine the amount of titrant delivered. Use
this, the concentration of the titrant, and the
stoichiometry of the titration reaction to
calculate the number of moles of reactant in
your analyte solution.
Example: It is found that 10.00 cm3 of 0.200 mol/dm3 aqueous
sodium carbonate requires 25.00 cm3 of hydrochloric acid to just
neutralize it. What is the concentration of the hydrochloric acid?
1. Calculate the amount in the solution of known concentration
Amount of sodium carbonate: 0.200 mol/dm3 x 0.0100 dm3 = 0.00200 mol
2. Write the balanced equation
Na2CO3 +
2HCl
2NaCl + CO2 + H2O
1 mol
2 mol
0.00200 mol
2x0.00200 mol= 0.00400 mol
3. Calculate the concentration of the unknown solution
[HCl] = 0.00400 mol/0.02500 dm3 = 0.160 mol/dm3
Titration with a pH meter
Titration with a pH meter follows the
same procedure as a titration with an
indicator, except that the endpoint is detected
by a rapid change in pH, rather than the color
change of an indicator.
Arrange the sample, stirrer, buret, and pH
meter electrode so that you can read the pH and
operate the buret with ease.
To detect the endpoint
accurately, record pH vs. volume of titrant
added and plot the titration curve(graph
of pH vs volume of titrant added.
Titration curves for strong acid – strong base titrations
LIGHT ABSORPTION
• If a reaction produces a precipitate, then the
time taken for the precipitate to obscure a mark
on a pieces of paper under the recation vessel
can be used as a measure of reaction rate.
This is suitable for the reaction:
S2O32-(aq) + 2H +(aq)
H2O(l) + S(s)
which produces finely divided precipitate of
sulfur.
• If the reaction involves a colored reactant
or product, then the intensity of the color
can be used to monitor the concentration
of that species.
• This can be done by comparing the color
by eye againist a set of standard solutions.
• It is better to use an instrument that
measures the absorbance.
Absorbance and Beer Lambert’s Law
Because light absorption is a function
of the concentration of the absorbing
molecules, a more precise way of
reporting intensity of absorption is by
use of the Beer-Lambert Law:
Absorbance = -log(I1 / I0) = acl
where:
a = molar absorptivity coefficient(how
strongly a compound absorbs light)
c = molar concentration of solute
l = length of sample cell (cm)
I0: initial intensity
I1: Final intensity
Absorbance is found directly by UV/visible
spectrophotometers.
CLOCK TECHNIQUES
• In some reactions the product can be
consumed with addition of another
substance. When all of this substance is
consumed an observable change will
occur. The time taken for this corresponds
to the time for a certain amount of product
to have been formed. So rate α 1/time.
• The classical reaction for this technique is:
H2O2(l) + 2H+(aq) + 2I-(aq)
2S2O32-(aq) + I2(aq)
2H2O(l) + I2(aq)
S4O62-(aq) + 2I-(aq)
The blue color of iodine-starch complex
suddenly appears when all of the thiosulfate
has been consumed.