Bloomington in particular, the local soil contains CaCO 3, which donates a carbonate (CO 2-3 ) ion to help
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1 Introduction Blake Dircksen Standardization of HCl and Buffer Capacity Determination of Local Water Samples Acid rain has been a problem around the world since the Industrial Revolution. The ph of natural rain hovers around 5.6 and anything below a ph reading of 5.6 is classified as acid rain. This reading of ph = 5.6 is lower than that of tap water (ph = 7) due to solubility of carbon dioxide gas in the atmosphere. Acid rain comes about by the emissions of nitrogen and sulfur oxides which arise by high temperature combustions (nitrogen) or from burning coal containing sulfur. Those emissions into the air eventually react with the atmospheric CO 2 in a series of reactions that results in sulfuric acid (H 2 SO 4 ) and nitric acid (HNO 3 )(Robinson). This acidic water then falls to the earth which can have negative environmental effects such as acidification of water bodies, damage of trees at high elevations, and decay of building materials and paints (EPA). Luckily, soil and rocks around the world have a natural buffering capacity. A buffer will resist a change of ph by adding either weak acid or its conjugate base to the equilibrium. In Indiana and Bloomington in particular, the local soil contains CaCO 3, which donates a carbonate (CO 2-3 ) ion to help resist acidification. The purpose of this lab was to determine the efficiency at which local water samples, Clear Creek and Jordan River, can buffer an acidic solution. The acidic solution being used in this lab is a 0.01 M solution of HCl. Due to the inability to buy 100% pure reagents and the possibility of a fluctuating concentration of HCl, a standardization was done in week 1 to ensure the exact concentration of HCl. This reagent and its measured concentration was used for experiments and calculations in week 2 (Robinson 66).
2 Methods Week 1 was used to prepare a standardized HCl solution that could be utilized for week 2. The process of standardizing involves determining its concentration by titrating with a basic primary standard of a known concentration (Robinson 65). A 500 ml 0.01 M HCl solution was made from using a 0.1 M HCl stock solution. The correct amount of 0.1 M HCl stock solution to add was found by equation 1: M 1 V 1 = M 2 V 2 (1 M)(V 1 ) = (0.01 M)(500mL) Eq. 1 V 1 = 5 ml About 250 ml of distilled H 2 O was added to a flask, and then the 5 ml of 0.1 M HCl solution, which was finally diluted to a mark of 500 ml. With this HCl solution ready, trials to determine the exact concentration could begin. The 0.01 M HCl solution was titrated with enough Na 2 CO 3 that would react with approximately 35 ml of the HCl solution. To determine how much Na 2 CO 3 was needed, a look at the stoichiometric equation was needed: 2HCl (aq) + Na 2 CO 3 (aq) à H 2 CO 3 (aq) + 2NaCl (aq) From this, the mass of Na 2 CO 3 needed to react with 35 ml of HCl could be determined using equation 2: 35 ml HCl!!!.!"!"#!"#!"""!"!!!!"#!"!!"!!!"#!"#!"#.!""!!"!!"!!!"#!"!!"! = g Na 2 CO 3 Eq. 2 Four samples of Na 2 CO 3 were measured out and recorded in Table 1. They were diluted with about 25 ml of distilled water and bromocresol green indicator was added. The titrant ( 0.01 M HCl) underwent a rough titration trial, where an approximate endpoint was found, and three careful titration trials after that. The volume of titrant used was observed for each trial and recorded in Table 1. The equation for determining the final Molarity of HCl can be found in equation 3 in the results section. Week 2 used the M HCl (Table 1) solution found from week 1 to determine the buffer capacity in Clear Creek and Jordan River in Bloomington, Indiana. Before trials could begin, the ph meter needed calibration from known ph solutions. The probe was immersed in buffer solutions of 4, 7, and 10
3 ph to set the calibration. With a calibrated ph probe, ph readings from rain water, distilled H 2 O at atmospheric CO 2 equilibrium, saturated CO 2 in H 2 O, and saturated CaCO 3 in H 2 O were observed and recorded in table 2. Using the ph probe, titration with the standardized HCl could now be used to determine buffering capacity. With a clean buret, M HCl was added to 100 ml of each water sample (Clear Creek / Jordan River) in 0.5 ml increments and the ph was recorded after each addition. These values were recorded and plotted in Figure 1 and 2. After multiple trials of slow changing ph, the increment level was adjusted to 1.0 ml. This was continued and ph was recorded until the ph hit 4.3. Results Table 1 below displays the data collected in week 1 when standardizing the solution of HCl. A rough titration followed by three titration trials were tested to find the volume of HCl titrated. Once the volume of HCl titrated (after the blank calibration) was determined, the molarity of the HCl solution could be calculated using equation 3 shown below: 2HCl (aq) + Na 2 CO 3 (aq) à H 2 CO 3 (aq) + 2NaCl (aq)!!"#!"!!"!!!"#!"#! Mass (g) Na 2 CO 3!"#.!""!!"!!"!!!"#!"!!"! = Molarity of HCl Eq. 3!"!#$"%&!"#$%& (!)!!"#!"!!"!!!"#!"#! Ex) g Na 2 CO 3!"#.!""!!"!!"!!!"#!"!!"! = M HCl.!"#$! Trial Mass Na 2 CO 3 (g) Volume of HCl Solution (ml) Volume of HCl Solution after blank correction (ml) Molarity of HCl rough Blank Mean St. Deviation Table 1: Data table for the standardization of titrant (0.01 M HCl).
4 A mean and standard deviation of the Molarity of HCl was taken and recorded in Table 1. The mean value of HCl Molarity was used in the following experiment. Week 2 of the experiment used the standardized HCl from week 1 to determine the buffer capacity of Clear Creek and Jordan River, two local water sources. Once the ph probe was calibrated, measurements of various water samples we found: rainwater had a ph reading of 5.69, distilled H 2 O at Atmospheric CO 2 Equilibrium was 6.15, saturated CO 2 in H 2 O was 4.04, and saturated CaCO 3 in H 2 O was 7.75, as listed in Table 2 below. Figure 1 and 2 below show the ph vs Volume of HCl relationships in the Clear Creek and Jordan River water samples respectively. Identity of Sample Rain water 5.69 Distilled H 2 O at Atmospheric CO 2 Equilibrium 6.15 Saturated CO 2 in H 2 O 4.04 Saturated CaCO 3 in H2O 7.75 ph Clear Creek 7.46 Jordan River 7.80 Table 2: ph recorded using various aqueous solutions. ph Buffer Capacity of Clear Creek Volume of Titrated HCl (ml) Figure 1: ph versus Volume of HCl titrant to determine a buffer capacity for Clear Creek
5 Buffer Capacity of Jordan River ph Volume of Titrated HCl (ml) Figure 2: ph versus Volume of HCl titrant to determine a buffer capacity for Jordan River From the information gathered in Figure 1&2, a buffer capacity can be calculated for each sample. The buffer capacity expresses the mols of HCl required to bring 1 L of water to a ph reading of 4.3. The equation for buffer capacity can be found in equation 4 below: ml HCl Titrated!!!.!!"#$!"#!"#!"""!"!!!!.!"!!"#$%!"#$%&!"""!!"#!!"# = mmol HCl/L Eq. 4 Ex) 43.5 ml HCl!!!.!!"#$!"#!"#!"""!"!!!!!"#!"""!.!!!"#$%!"##$!"#$%&!!"# = 4 mmol HCl/L Clear Creek Environmental Sample mmol/l of HCL Clear Creek 4 Jordan River 3.5 Table 3: Buffer capacity readings for Clear Creek and Jordan river. A reading of mmol/l is a common buffer capacity unit. Typical values for buffer capacity are 0.05 to 2.0 mmol/l Discussion CO 2 in the air and anions in the soil are known for their acidic or basic properties they can have on water. This is why natural rain water has a lower ph than that of natural water (ph =7) the carbon
6 dioxide in the atmosphere dissolves in rain water to produce the weak acid, carbonic acid (EPA). This relationship can be seen below: CO 2 (g) + H 2 O (l) à H 2 CO 3 (aq) The rain water in Indiana is acidic (ph 5.6) because of the emissions put into the air. In 2011, Indiana produced 371,692 tons of SO 2 and 117,825 tons of NO x. This is a large number is due to the large amount of coal Indiana burns. A state like Colorado, which doesn t burn as much coal, produces only 43,320 tons of SO 2 and 48,962 tons of NO x (EPA). A saturated CO 2 aqueous solution would have a lower ph than that of a CO 2 aqueous solution at equilibrium because the saturated solution has a higher concentration of CO 2 dissolved in the water which would produce more carbonic acid, resulting in a lower ph reading. The ph readings for saturated CO 2 and CO 2 at equilibrium were 4.04 and 6.15, respectively. Alternatively, the saturated CaCO 3 water sample has a higher reading of 7.75 ph because when CaCO 3 dissociates in H 2 O, its carbonate anion (CO 2-3 ) reacts with the hydrogens in water, resulting in a higher ph reading. These ph readings by the probe reflect the relationships that should be expected in nature. Because Bloomington, Indiana contains limestone and therefore calcium carbonate (CaCO 3 ), it is expected that the local water samples have a high buffering capacity. The buffer capacity again refers to the quantity of acid or base that 1.0 L of the buffer can accommodate without undergoing a significant ph change. Since the water samples have buffering capabilities, it would be expected for the slope of the titration curve to be slowly changing. In a solution of unbuffered water, the slope for the titration curve would look something like this:
7 Buffer Capacity of Water ph vol. of 0.1 M HCl (ml) Figure 3: Buffer capacity of water, an unbuffered solution (Robinson). where a small amount of HCl can lead to a big drop in ph. From looking at Figure 1 Clear Creek data, it can be observed that it takes a relatively large amount of HCl, approximately 43 ml, to bring the ph reading down to 4.3. When comparing the data from the two local water samples, Clear Creek and Jordan River, it can be determined that Clear Creek has a higher buffer capacity of 4 mmol/l of HCl while Jordan river had only 3.5 mmol/l HCl. These values for the local water samples show a high capability for buffering acids. Typical values for buffering capacity range from 0.5 mmol/l to 2.0 mmol/l, so these readings are rather high. It is possible that these high readings are the naturally occurring alkalinity measures (buffer capacity) of the local water samples, and that since Bloomington has such a high concentration of calcium carbonate, these readings are accurate. However, it is also possible that there could have been sources of error in this experiment. Errors in the titration process could have led to inaccurate readings. At times, 1 to 2 ml of M HCl was added, causing a slight overshoot in the ph readings. In the standardization of HCl in week 1, it was also difficult to determine a uniform green color of indicator, and more trials would have been beneficial to ensure accurate findings.
8 References "Effects of Acid Rain - Surface Waters and Aquatic Animals." EPA. Environmental Protection Agency, 08 June Web. 20 Sept < Robinson, J.K., (2010). C118 Laboratory Manual, Principles of Chemistry and Biochemistry II, Indiana University. Hayden McNeil Publishing, Plymouth, MI, pg
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