Investigation of the effect of high te t e m p e r a t u r e e x p o s u r e o n t h e o x i d i z i n g power of sodium hypochlorite bleach
AUTHOR
Kai Roger Smith Hove
SUPERVISOR
Anniken T. T. Øverås
SCHOOL
Arendal Videregående Skole
COMPLETED
05.11.2011
WORD COUNT
3978 words
Abstract Liquid bleach, in the form of sodium hypochlorite (NaClO), is an unstable compound, and is subject to a number of decomposition pathways depending on pH, UV radiation, transition metal contamination and temperature. The scope of this experiment includes an investigation of the effect of the latter: “How does exposure to high temperatures affect the oxidizing power of sodium hypochlorite bleach (on iodide)?” A hypothesis predicting reduced oxidizing power at higher temperatures was made on the grounds of a proposed thermal decomposition pathway, which predicts the formation of chlorate and chloride salts. A high concentration of the chlorate ion, which displays a weak standard reduction potential, would reduce the oxidizing power of the solution. In the experiment, four undiluted samples of commercial bleach solution were heated within a temperature range of 30-40 °C, 50-60 °C, 70-80 °C and 90-100 °C respectively, over a period of 2 hours, with a reflux condenser. After allowing cooling overnight, a series of iodometric titrations were performed to determine the oxidizing power of the solutions. This involved the oxidation of excess iodide to iodine by hypochlorite in acidic media. The liberated iodine was then reduced back to iodide by titration with sodium thiosulfate with a starch indicator. The oxidizing power of the bleach was subsequently extrapolated with respect to how much iodide that was oxidized under each temperature condition. Unfortunately, no strongly justified conclusions could be drawn due to the significant spread in the data, which indicated the presence of a systematic error. However, qualitative data in the form of crystallization of a hygroscopic substance was observed in the two samples heated above 70 °C, something which partly supports the initially proposed chlorate-forming decomposition pathway. Words: 294
Table of contents
Section
Title
Page
Abstract 1
Introduction and research question
1
2
Background theory
2
3
Hypothesis
3
4
Method
4
4.1
Heating the bleach
6
4.2
Iodometry
7
5
Results
8
6
Analysis
10
7
Conclusion
11
8
Evaluation
14
9
8.1
Error sources and improvements in experimental design
12
8.2
Literature evaluation
13
Unanswered questions and further research Bibliography
13
1 Introduction and research question Sodium hypochlorite (NaClO), was first synthesized in the late 1800s through the efforts of French chemists Claude Louis Berthollet and Antoine Labarraque. It came into widespread use in the 1930s, and to this day it remains a high-demand product. Its desirable oxidizing and antibacterial properties lend themselves to a wide range of applications, from many industrial and professional uses such as in endodontics (Clarkson, 1998), to general household cleaning and bleaching of clothes. However, the use of sodium hypochlorite comes with an inherent drawback; it is an unstable compound, and is subject to degredation through various decomposition pathways catalyzed by factors such as UV radiation from sunlight, transition metal contamination and temperature. (Harms et. al., 2010) The scope of this investigation includes the effect of the latter. For both producers and end users of bleach alike, it is important knowing to what extent high temperature exposure affects the oxidizing power of the bleach solution, and in light of this knowledge, handle the solution correctly with respect to its production, storage and use in order to maximize its efficiency. The oxidizing power of the bleach solution can be inferred from to the extent to which it can successfully oxidize other compounds. In this investigation, iodide has been used as a reducing agent.
This leads to the research question of this essay: “How does exposure to high temperatures affect the oxidizing power of sodium hypochlorite bleach on iodide?” During preliminary investigation of the subject, it would appear that previous research on the effect of temperature on sodium hypochlorite solutions has yielded contradictory results, something that was also noted by Frais et. al. (2001). While some researchers have reported little to no effect of high temperatures on the oxidizing efficacy of sodium hypochlorite solutions (Gambarini, G., De Luca, M. & Gerosa, R., 1998), Frais et. al. reported a significant loss in oxidizing efficacy. This unresolved area of research was something that peaked my personal interest, and why I deemed it worthy of further investigation.
1
2 Background theory
Commercial bleach solution is produced by bubbling chlorine gas through an aqueous solution of excess sodium hydroxide. The resulting solution of bleach consists of sodium hypochlorite, sodium chloride, water, as well as excess sodium hydroxide. (Harms et. al., 2010) Cl2 (g) + 2 NaOH (aq) ! NaClO (aq) + NaCl (aq) + H2O (l) In the precense of water, the sodium hypochlorite will then dissociate into ions, NaClO (aq) ! Na+ (aq) + ClO – (aq) whereas the hypochlorite ion ClO – is responsible for the oxidizing action of the bleach, while the sodium is simply a spectator ion. As previously mentioned, the problem with liquid sodium hypochlorite in solution is that it has a tendency to decompose through reactions with itself. Sodium hypochlorite is subject to several decomposition pathways, of which the primary is a reaction where three moles of hypochlorite ions react to form one mole of chlorate and two of chloride. (Zoller, 2009) This is a disproportionation reaction, as Cl is simultaneously oxidized and reduced. 3 ClO – (aq) ! 2 Cl – (aq) + ClO3 – (aq) Another, much less prevalent way of degradation is the oxygen-producing decomposition pathway, which results in the formation of sodium chloride and oxygen gas. (Lister, 1956) 2 OCl – (aq) ! 2 Cl – (aq) + O2 (g) This occurs to a negligable extent and has not been proven to be temperature-dependent, but the presence of certain transition metal ions, namely copper, nickel, iron, cobalt and manganese, have been shown to greatly accelerate this reaction through catalytic activity. (Lister, 1956) Normally, under proper bleach production conditions, the amount of these ions would be so insignificant as to not catalyze any significant decomposition. For that reason, the discussion in this essay will take place under the assumption that chlorate formation is the main decomposition pathway. The problem that arises is the conflicting presences of the hypochlorite ion, ClO –, versus the chlorate ion, ClO3 –. When we are discussing in terms of the oxidizing efficacy of these ions in bleach solution, which is the more desirable oxidant? There appears to be conflicting theory surrounding this problem: ClO3 – is a highly oxidized anion, displaying chlorine in the +5 oxidation state, versus ClO –, which is +1. As as ClO3 – displays a higher oxidation state, it follows that it should also be expected to show stronger oxidizing properties. However, this does not seem to be the case.
2
The following table, adapted from Cotton et. al. (1988), shows the standard reduction potential in volts for the successive oxyacids of chlorine through their reduction in acidic media. Each species has its own standard reduction potential, which is a measure of the species’ affinity for electrons - with greater potential, it shows a greater tendency to be reduced. Table 1; Standard electrode potentions in volts for the reduction of successive oxyacids of chlorine in acidic media Reduction half-reaction
E ° (V)
H+ + HClO + e – ! #Cl2 + H2O
+1.63
3H+ + HClO2 + 3e – ! #Cl2 + 2H2O
+1.64
6H+ + ClO3 – + 5e – ! #Cl2 + 3H2O
+1.47
8H+ + ClO4 – + 7e – ! #Cl2 + 4H2O
+1.42
From this data, we can see that the hypochlorite ion displays the strongest tendency to be reduced, while this trend decreases with the successive oxychlorines (with the exception of the chlorite ion, ClO2 –, which is a marginally stronger oxidizer than its predecessor). This trend of weaker oxidizing power in relation to higher oxygen content seems paradoxical, considering it conflicts with the usual trend of higher oxidizing potential in relation to higher oxidation states. One possible explanation for this discrepancy relates to the higher ions’ increased kinetic stability. Chlorate’s relative ineffectiveness as an oxidizer compared to hypochlorite could stem from the increased number of electrons involved in the formation of " -bonds in the transition from ClO – to ClO3 –. (Tyagi, 2009) This, combined with the shielding of the central chlorine atom by the increased number of oxygens, would add to the ion’s stability, making it more unreactive. The Lewis structures of the ions are illustrated below.
Figure 1 Lewis structures of hypochlorite and chlorate
Hypochlorite
Chlorate
-bonds
!
3 Hypothesis My hypothesis is that as higher temperatures will result in an increase in kinetic energy, this should increase the rate of the reaction where hypochlorite decomposes to chlorate. Chlorate being the less effective oxidizer, as demonstrated by its lower reduction potential and more stable structure - should oxidize less iodide. In other words, I expect heated bleach to be less effective.
3
4 Method
The first step of this experiment involves the heating of the bleach. As perfectly accurate temperature control is not possible with the available equipment, four ranges of temperature will be employed instead: 30-40 °C, 50-60 °C, 70-80 °C and 90-100 °C. Furthermore, high temperatures up to 100 °C was chosen because sodium hypochlorite decomposition is usually a very slow process that happens over several months time under room temperature conditions. Thus, higher temperatures should be able to accelerate decomposition enough to produce quantitiatively significant results within the time constraints of the investigation. A heating time of 2 hours was employed. The method of heating was partially adapted from the method by Frais. et. al. (2001). As opposed to their method, which consisted of heating the bleach uncovered, I will be utilizing a reflux condenser to retain the concentration of the solutions, which is an important control variable in this experiment. As previously mentioned, decomposition of hypochlorite largely results in the formation of chlorate ions. In light of this, a very accurate method of measuring the extent of bleach decomposition could be through the detection of chlorate ions by ion chromatography or NMR spectra. Consequently, the oxidizing power could be extrapolated from the chlorate ion concentration, which should be inversely related. However, while this method would provide us with very accurate results, such sophisticated equipment is not readily available. Instead, a simpler, though admittedly less accurate method can be used to indicate the oxidizing power of the solutions $ iodometry. In acidic solution, hypochlorite ions in the bleach will oxidize iodide ions to form iodine, chloride ions and water. ClO – (aq) + 2 I – (aq) + 2 H+ (aq) ! I2 (aq) + Cl – (aq) + H2O (l) The liberated iodine forms an equilibrium with itself, giving rise to triiodide ions, which turn the solution distinctly yellow: I2 (aq) + I$ (aq) I3$ (aq) The solution can then be titrated with sodium thiosulfate. The thiosulfate ion reduces the iodine back to iodide, gradually draining the solution of its color in the process. 2 S2O32– (aq) + I2 (aq) 2 I$ (aq) + S4O62– (aq) With the solution gradually taking on only a faint yellow color, the endpoint can be difficult to determine. For that reason, starch solution is added. The iodine forms a reversible, deeply dark-blue colored complex with the starch that is useful for accurate endpoint detection. When the dark blue color disappears, this indicates that all the iodine has been reduced back to iodide ions. From the molarity of the titrant added and the reaction stoichiometry, it is possible to deduce how many moles of iodide were oxidized by the hypochlorite. This would provide an indication of the oxidizing power of the solution. In this experiment I expect thermally decomposed bleach would contain higher chlorate content. Chlorate, being the weaker oxidant, should oxidize less iodide ions.
4
Table 2; Variables
Variable type
Variable
Method of measurement / control
Oxidizing power
Indicated by amount of iodine liberated through oxidation, determined through iodometric titration
Temperature
Record using a temperature node attached to a digital device and lowered into solution. Temperature controlled through adjusting the electric heating mantles.
Type of bleach
Klorin, produced by Lilleborg A/S.
Volume of bleach
50.0 mL will be heated.
Concentration of bleach
Utilize reflux condenser in order to recondense vapors when heating. Extracted samples will be diluted by the same amount before titration.
Heating time
All four samples will be heated for 2 hours.
Na2S2O3 and KI solutions
Na2S2O3, the titrant, can decompose over time, and KI can be gradually oxidized by air, volatilizing the iodine. Therefore, these solutions will be prepared from solids before use.
Addition of starch
Must happen close to the end point, otherwise the iodine can form an irreversible complex with the starch.
Dependent
Inependent
Control
5
4.1 Heating the bleach
For the process of heating, four round-bottom boiling flasks were each filled with 50.0 mL of undiluted bleach, measured out in a measuring cylinder. A reflux condenser was utilized in order to prevent evaporation. The flask and tube was put on top of an electric heating mantle and secured with a clamp. The setup was placed in the fume hood, as to make away with what small amount of vapors that may escape. A temperature node attached to a digital PASPORT Xplorer GLX device was lowered into the flask, and porcelain boiling chips were added to prevent sudden, uncontrolled boiling. The heat was then turned on for all flasks, and the time was recorded from the point when all the solutions were in their designated temperature range, which took roughly 8 minutes. The solutions were heated for 2 hours before the heat was turned off. Interestingly, hints of a crystalline precipitate had formed in the solutions heated above 70 °C. The samples were left in the fume hood overnight to restabilize to ambient temperature and for vapors to recondense.
Figure 2 Apparatus setup
Electric heating mantle
6
4.2 Iodometry
The following day, the bleach solutions were dismounted and remeasured and were found to have lost <1 mL of water, something which should not impart a significant influence on the results. 4.96 g of sodium thiosulfate pentahydrate was dissolved in 1.00 L of distilled water to make a solution of 0.0200 M Na2S2O3. 1.66 g of potassium iodide was dissolved in 100 mL of water to make a solution of 0.100 M KI. A 2.00 g/L starch solution was prepared under heating until complete dissolution. 10.0 mL of each of the bleach solutions was pipetted into to a volumetric flask, diluted to 250 mL with distilled water and homogenized. Followingly, a 5.00 mL aliquot of the dilute solution was pipetted into an Erlenmeyer flask and diluted with another 20.0 mL of water, measured out in a 25.0 mL measuring cylinder, to increase the volume. 2.00 mL of a 6.00 M stock solution of HCl was added to the titration to flask to provide the acidic environment needed for the oxidation reaction to proceed. After this, 2.00 mL of the KI solution was added, initiating the oxidation reaction between ClO – and l –. Rapidly after, the solution was titrated against Na2S2O3 from the burette until the solution was only faintly yellow, after which 2.00 mL of the starch solution was added, turning the solution strongly blue. Titration was continued with consistent swirling of the flask until the solution just turned colorless, indicating complete reduction of the iodide.
7
5 Results
Quantitative raw data Table 3; Volume of Na2S2O3 required to reach endpoint in iodometric titration Vol. Na2S2O3 in mL ± 0.05 30-40 °C
50-60 °C
70-80 °C
90-100 °C
Control (RT)
Trial 1
12.40
8.90
8.90
5.85
7.40
Trial 2
11.00
9.30
10.05
5.40
5.50
Trial 3
11.05
8.95
11.35
8.55
6.10
Qualitative raw data After one hour of heating, a white, crystalline substance had accumulated at the bottom of the flasks for temperature ranges 70-80 °C and 90-100 °C. After the powder was extracted from another sample and left to dry in air, it started visibly turning into a greyish slurry. As for the iodometry; after adding KI to the bleach solution, it turned yellow. Following titration with Na2S2O3, the color faded until it was only faintly yellow - at which point starch was added, and the solution turned black-blue. Titration was continued until it turned colorless. Several minutes after the titration, the solution had again taken on a very faint blue tint.
8
Processed data The following table shows the calculated number of mols for each reaction. The number of mols is the amount of sodium thiosulfate required to completely reduce the iodine that resulted from the hypochlorite oxidation of iodide. To reduce random error, the mole calculations were based on an average of three parallell trials for each temperature range.
Table 4; Processed data Vol. Na2S2O3 in mL ± 0.05 30-40 °C
50-60 °C
70-80 °C
90-100 °C
Control
Trial 1
12.40
8.90
8.90
5.85
7.40
Trial 2
11.00
9.30
10.05
5.40
5.50
Trial 3
11.05
8.95
11.35
8.55
6.10
1.40
0.40
2.45
3.15
1.90
11.50
9.10
10.10
6.60
6.30
1.840 ! 10 "#
2.202 ! 10 "#
1.320 ! 10 "#
1.260 ! 10 "#
Spread Average
2.300 ! 10 "#
mols Na2S2O3 / I-
The number of mols of oxidized iodide was calculated followingly, based on the known volume and concentration of the sodium thiosulfate solution and the stoichiometry of the reaction equations.
Example calculation n(Na2S2O3) = cV = 0.0200 mol dm%& x 0.0115 dm& n(Na2S2O3) = n(I%) = 2.300 ! 10 "# mol
The results are presented in the following diagram: Figure 3 Mols Ioxidized
3.000 ! 10 "# 2.700 ! 10 "# 2.400 ! 10 "# 2.100 ! 10 "# 1.800 ! 10 "# 1.500 ! 10 "# 1.200 ! 10 "# 9.000 ! 10 "$ 6.000 ! 10 "$ 3.000 ! 10 "%
Control (RT*)
30-40
50-60
* room temperature
9
70-80
90-100
Temperature range in °C
6 Analysis We would have expected the solutions that were exposed to higher temperatures to have undergone a faster rate of decomposition to chlorates and chlorides, thus reducing the amount of hypochlorite available for oxidation, resulting in less oxidized iodide, and ultimately requiring less volumes of titrant to completely reduce the oxidation products. Disappointingly, it appears that the experiment was not a success. The data cannot be interpreted with respect to the hypothesis, as the results do not show a consistent trend, and a significant spread can be observed even between parallel trials. This suggests the precense of a systematic error. There are two major discrepansies that discredit the recorded data: i) there is a large spread across most parallell trials, the largest being a difference of a whole 3.15 mL of titrant in the 90-100 °C sample, and ii) the control sample, stored at ambient conditions, displays the lowest oxidizing power of all the solutions, which is the opposite of what is expected. Some parallel trials show a gradual positive increase in the titrant added (90-100 °C), some a negative one (30-40 °C), and one remained relatively consistent across parallel trials (50-60 °C). They are certainly not uniform as they should be - preferably, titrations should agree to the nearest drop. It seems the direction of the error goes both ways, suggesting that the cause of the error present is of a fluctuating nature. For that reason it is unlikely that the error is related to the standardized thiosulfate solution - if there was a problem relating to its quality the error should have been in a uniform direction. Other external factors must have played a role, which are discussed under section 8.
10
7 Conclusion In light of the inconsistensies of the quantitative data, giving a definite answer is difficult, as they do not display any consistent trends. In all likelihood, this was caused by a yet unidentified systematic error. As such, we cannot make a justified assertion as to whether exposure to high temperatures affects the oxidizing power of sodium hypochlorite bleach, and the hypothesis cannot be accepted. However, some qualitative data supports the hypothesis. While the qualitative data is not sufficient in itself, it can be taken into consideration and contexualized in relation to other research. In the bleach samples of 70 °C and above, visible salt crystallization had occured $ most likely a mixture of sodium chlorate and sodium chloride, as per the proposed main decomposition pathway. This suggests the solution was sufficiently decomposed as to be saturated with ions. A significant precense of NaClO3 could be supported by the observation that the salt was highly hygroscopic, as it turned into a slurry over time. This is a property that characterizes NaClO3. (Eagleson, 1988) This suggests a considerable precense of chlorate ions, but it remains unknown whether this correlates with the solution’s oxidizing power. The data is partially supportive of the results found by Frais et. al. (2001). With a storage temperature of 37 °C during the course of 6 months, their bleach sample was found to have experienced a 62% reduction in chlorine species available for oxidizing action. Solid salt formation had also occured in their sample, which is also attributed to be a mixture of sodium chlorate and sodium chloride, but this was not experimentally confirmed. With more parallell trials and the implementation of the improvements discussed under section 8, it should be possible to make valid observations in this experiment.
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8 Evaluation 8.1 Error sources and improvements in experimental design Seeing as errors appeared to occur in a random direction, this makes it is difficult to pinpoint one defining source of error. A list of errors and their potential improvements will be discussed. A major weakness of the experiment was that it was only performed once. Therefore, regardless of the direction of the data, there may not be strong enough support to claim that the results show a consistent pattern. A repetition of the procedure at least three times could provide grounds for stronger evidence. Also, doing more than three parallell runs for each titration, preferably around ten, would greatly help reduce random error. As for the iodometry, a possible source of systematic error includes oxidation of the iodide ions by O2 in the air. Considering the relatively short time taken to complete a titration it may not exert a big difference (weak reoxidation was only visibly apparent 10-15 minutes after reduction with thiosulfate), but it is a probable error source that should be limited nonetheless. One way of doing this could be by filling the titration flask with a less reactive gas, like N2 or CO2, during the titration, something which would displace the oxygen and limit the influence of O2 oxidation. Another area of improvement is the temperature control under heating. The flasks with solutions were heated directly in electric heating mantles, which proved to not result in very precise temperature control. Within very small variations on the heat setting, temperature was prone to fluctuate within a spread of 10 °C. For that reason, temperature ranges had to be chosen as opposed to specific temperatures. This limits the precision of the data. With more precise temperature control, which could be accomplished through slower, more deliberate heating over a longer period of time, this would reduce the influence of random error and allow for more justified statements regarding the effect of a particular temperature on the solution. An extensive treatment of quantitative uncertainties in measurements was omitted. This could be improved by taking into account all uncertainties and calculating the percentage uncertainty. An inherent limitation of the procedure is that we cannot ascribe the oxidizing action to a specific ion - we can only draw conclusions based on the theories of the ions’ structures and reduction potentials. The oxyanions of chlorine display similiar chemical properties and are prone to rapid interconversion, making it difficult to isolate or distinguish between them by chemical tests. This limitation could be overcome by using more sophisticated techniques; the relative presences of certain anions is frequently measured using ion chromatography. With this method, the samples could be double checked to see if the hypochlorite anion concentration does in fact correlate with the solution’s oxidizing power. Another method could involve the use of nuclear magnetic resonance spectra, calibrated for detecting the presence of the different ions.
12
8.2 Literature evaluation The works cited in this investigation are taken from, through my best judgment, credible authors under established publishers, academic journals and reviews from industrial research institutions. Therefore, I assume with fair certainty that the essential literature consulted is accurate.
9 Unanswered questions and further research Some unanswered questions remaining are: • • • •
What is the ratio of chloride to chlorate ions in the solution before and after heating? What is the oxidation-reduction potential of the solution before and after heating? (This could be determined using an ORP meter.) Does the duration and intensity of heating affect how the bleach decomposes? Short and intense exposure to heat or long exposure at a mild temperature? Can bleach solutions be effectively preserved at cold temperatures over longer time periods?
The experiment leaves multiple opportunities for further research. Since we are discussing temperature, one could also approach an investigation from a kinetic viewpoint: • •
What is the rate and reaction order of the decomposition process? Does the reaction satisfy the rule of thumb for a doubling in rate for every 10 °C increase in temperature? (Brown, C. & Ford, M., 2009)
Other possible areas of research include the reaction mechanism for transition metal catalyzed decomposition, and the effect of UV radiation (in the form of sunlight exposure, for instance). This combined knowledge should provide a good foundation for understanding how sodium hypochlorite bleach should be produced and stored in order to maximize its oxidizing efficacy.
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•
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