Environmental Chemistry U6220
Lab #3
Part 1: Investigation of the Carbonate-Bicarbonate System (Alkalinity of Aqueous
Systems)
Introduction:
An aqueous solution of carbon dioxide produces a mixture of carbonate and bicarbonate
ions. Determining the carbonate and bicarbonate ions in each other's presence is often
important in environmental chemistry.
a) CO2(g) + H2O(l)  H2CO3 (aq)
b) H2CO3 (aq)  HCO3-(aq) + H+(aq)
c) HCO3-(aq)  H+(aq) + CO32-(aq)
The alkalinity of water is its capacity to act as a base by reacting with protons. There
exists a fundamental difference between the expression of acid-base properties of pH and
alkalinity. Whereas the pH can be considered to be an intensity factor which measures the
concentration of alkali or acids immediately available for reaction, the alkalinity is a
capacity factor which is a measure of the ability of water sample to sustain reaction with
added acids (in a sense, it is the ability of a water body to neutralize added acids). In
practice, it may be determined by measuring the number of moles of H+ required to titrate
one liter of a water sample to the end point. Acidification of a lake in its natural setting is
itself analogous to a macro-scale titration and lakes are sometimes termed well-buffered,
transitional, or acidic depending on their position on the titration curve.
Alkalinity is therefore a useful measure of the capacity of water to resist acidification
from acid addition (e.g. acid precipitation). The presence of carbonate, bicarbonate, and
hydroxide ions imparts the alkalinity of natural or treated waters. Initially, your water
samples may contain bases and will contain a positive alkalinity.
Question 1. Based on the pH measured in you initial samples (seawater and mangrove
swamp water):
a) What contributes to the pH differences you observe in your two samples (i.e. in
other words, what is the origin of the [H+] and alkalinity you observe).
b) What do you think is the most important carbonate species present in you
samples, and,
c) Is [OH-] an important contributor to the total alkalinity? Why? How should you
incorporate this base in any calculations associated with this experiment?
The addition of acid to the water sample will convert the carbonate to bicarbonate
(reverse of reaction c) until no carbonate remains. The addition of further acid will
convert the bicarbonate to carbonic acid until no more bicarbonate remains (reverse of
reaction b). The carbonate and carbonic acid equivalence points may be determined
either by titration using indicators or by pH titration.
The first end point determined (in the pH range 8.3-10) represents the completion
(equivalence point or stoichiometric end point) of the following reaction:
H+(aq) + CO32-(aq)  HCO3-(aq)
i.e. the carbonate has been neutralized by the acid-forming bicarbonate ions.
In the pH range 3.2-4.5, all of the bicarbonate ions initially present in the water sample,
together with all of those produced from the reaction of the carbonate ions, will be
neutralized. The resulting alkalinity is known as the total alkalinity.
HCO3-(aq) + H+(aq)  CO2(g) + H2O(l)
Question 2. Please write the equation for the estimated total alkalinity (what are the
assumptions you need to take? Are these valid? Say why)
The importance of the carbonate/bicarbonate system in natural waters stems from their
ability to act as buffers in natural waters. The oceans are described as being buffered
since relatively large quantities of acid or base can be added to seawater without causing
much change to its pH. However, many freshwater lakes do not have a large buffer
capacity and consequently a small addition of acid (e.g. from acid precipitation or
industrial effluent) can cause large changes in pH without warning. The carbonate
alkalinity and the total alkalinity are useful for the calculations of chemical dosages
required in the treatment of natural water supplies.
Summary of the Method:
Alkalinity is measured by titrating a water sample with sulfuric acid. The Vernier sensor
is used to monitor pH during the titration. The equivalence point will be at a pH of
approximately 4.5, but will vary slightly, depending on the chemical composition of the
water. The volume of the sulfuric acid added at the equivalence point of the titration is
then used to calculate the alkalinity of the water.
Material checklist:
1) Computer
2) Vernier computer interface
3) Logger Pro
4) Vernier pH sensors (2)
5) Sampling bottles
6) 25-50 ml burets
7) 100 ml graduated cylinder
8) 250 ml beakers (2)
9) Wash bottle with DI water
10) Utility clamps & Ring stand
11) 0.01 M H2SO4 solution
12) Magnetic stirrer and bar
Procedure:
pH Titration
Note (Caution!): Please wear gloves and safety goggles to perform this experiment and
beware that H2SO4 is corrosive. Avoid spilling it on your skin or clothing
1) Obtain at least 300 ml of seawater into a beaker (seawater and mangrove swamp
water).
2) Using a 100 ml graduated cylinder, carefully add 100 ml of each water sample in each
beaker
3) Place each beaker on the base of a magnetic stirrer and drop a stir bar carefully into
the beaker. Set the stirrer to a speed that mixes the sample well, but does not splash.
4) Keep the water away from the computer at all times.
5) Prepare the computer for data collection by opening “Test 11 Alkalinity” from the
Water Quality with Computers experiments files of Logger Pro. On the Graph
window, the vertical axis has pH scaled from 2 to 10 units of pH. The horizontal axis
has volume scaled from 0 to 20 ml. There is also a Meter window which displays real
time pH readings.
6) Insert the electrodes of the pH meter into the beaker.
7) Ensure complete coverage of the electrodes. It is essential that adequate clearance is
achieved between the electrodes and the magnetic stirrer or the stir bar will not rotate.
8) Place the burette (previously filled with 0.010M sulfuric acid) over the apparatus so
that acid can be run slowly into the beaker. Ensure that you have sufficient room to
turn the tap of the burette freely.
9) You are now ready to perform the titration. This process goes faster if one person
manipulates the buret while another person operates the computer and enters
volumes.
10) Click Collect to start data collection
11) Monitor the pH value on the computer screen. Once it has stabilized, click Keep.
12) Type 0 (the buret volume in ml) in the edit box. Then press Enter.
13) Add a small quantity of H2SO4 titrant (enough to lower the pH about 0.2 units). When
the pH stabilizes, click Keep.
14) Type the current buret reading to (the nearest 0.01 ml) in the edit box, then press
Enter.
15) Continue adding H2SO4 solution in increments that lower the pH by about 0.2 pH
units and enter the buret reading after each increment. When the graph shows the pH
value beginning to drop more quickly (at approximately 5.5). change to one-drop
increments. Enter a new buret reading after each addition. Note: It is important that
all additions of acid in this part of the titration be exactly on drop in size.
16) When the pH values start to flatten out (approximately pH 4), again add larger
increments that lower the pH by about 0.2 pH units, and enter the buret readings after
each increament.
17) Continue for two to three more additions, or until the graph clearly shows that the pH
has leveled off again.
18) Click Stop when you have finished.
19) Rinse the pH sensor with DI water from the wash bottle. Use a second beaker to catch
the rinse water. Return the sensor to the storage solution bottle and tighten the cap.
Question 3: Graph your results where pH appears on the y-axis and volume H2SO4
appears on the x-axis.
Question 4: (Calculations - Enter all your data in Table 1 below):
1) Determine the volume of H2SO4 titrant added at the equivalence point of the titration.
The equivalence point is the point where the titration curve makes the steepest drop in
pH.
a. Find the H2SO4 volume just before the start of the steepest pH change
b. Find the H2SO4 volume just after the start of the steepest pH change
c. Calculate the average of these points by adding them together and dividing
by two. Record this number, which represents the exact volume of H2SO4
added at the equivalence point (round to the nearest 0.1 ml).
(For an independent determination of the end points, plot the change in pH divided by the
change in volume on the y-axis, against the volume of 0.01M H2SO4 added, on the xaxis).
2) Calculate the mole of H2SO4 used to reach the equivalence point.
3) The reaction occurring in this titration is:
H2SO4 + CaCO3  CO2 + H2O + CaSO4
Based on the mole ratio of H2SO4 to CaCO3, calculate the moles of CaCO3 reacted at the
equivalence point
4) Calculate the mass in grams of CaCO3 in the sample. Convert to milligrams.
5) Calculate total alkalinity in mg/l of CO32-.
6) This value is equivalent to how many moles of [H+] neutralized?
Question 5: Repeat the whole procedure with mangrove swamp water.
Volume H2SO4 at
the equivalence
point (ml)
Sample
Alkalinity
(mg/l of CO32-)
moles of [H+]
neutralized
Seawater
Mangrove water
Part 2: Investigation of environmental parameters in diverse water systems
Introduction:
In this part of the lab exercise, you will be evaluating water quality parameters from
diverse types of water samples within and across different ecosystems.
Parameters measured:
- pH
- NH4+
- NO3- Dissolved Oxygen (DO)
Summary of the Method:
Measure all parameters for each sample using the previously set up Vernier sensors.
Record the value from the meter on the computer screen and enter this value in Table 2
below (make sure you rinse the sensor after each measurement and that you place it back
in its storage bottle).
Sample
Seawater
Mangrove waters:
- Upland fresh water
- Black Mangrove
- Red Mangrove
- Inlet to Algal Scrubbers
- Out of Algal Scrubbers
Algal scrubbers
Rain Forest:
- Mountain (stagnant)
- Waterfall (reservoir)
- Stream
- Soil percolate
pH
NH4+
NO3-
DO
Question 6: Explain the observed values in terms of the different environments they
represent within each ecosystem and across ecosystems. Place a particular attention to
these following environments:
- Black Mangroves
- Rain forest soil percolate
- Algal mats
a) In the black mangrove waters, can you explain the observed pH considering that these
waters should contain a relatively high alkalinity (brackish waters)
b) Are the rain forest soil percolate waters different from those found in surface waters?
How and why? What would that mean for example for transport of metals along a soil
profile?
c) In the algal mats, do you expect any parameter(s) to change over time? Which one(s)?
In which direction and by what magnitude? Over what time scale?