Full IS Paper (Marissa Walter)

AIES Independent StudyJune 4, 2015

Hydrolysis Kinetics and Efficiency for Hydrogen Production On-board Vehicles

Marissa WalterAdvisor: Tareq Abu-Hamed

Abstract:

Hydrolysis, or ‘water splitting’, emits no pollution, and recent studies suggest that solar energy could be used to recycle the metal oxidized in the process, which could close the loop of this process and make a storable and transportable, in the form of Hydrogen (H) fuel form of solar energy, that is more competitive with fossil fuels. This study investigated efficient hydrogen production, for use on-board vehicles, by measuring the effects of the catalyst, Sodium Hydroxide (NaOH), concentration on the rate of production of hydrogen gas. The time required to produce 50mL of H gas was recorded for exactly calculated amounts of Water, Aluminum foil, and NaOH. A positive correlation was seen between NaOH concentration and H production rate, and a significantly fastest rate for the concentration of 2.0M.

Introduction:

In recent decades, we have come to understand the serious consequences of fossil fuels, but energy is still critical for today’s higher standard of life. Global energy demand is currently 16.9TW and is expected to continue to grow with population. Oil resources and reserves are expected to last for 40 to 150 years (Katz, 2015), but to keep using oil as usual will increase emissions of CO2, methane and nitrous oxide - emissions, which are already unprecedented,1 and have deemed extremely likely to be a major cause of climate change (IPCC Fifth Assessment Synthesis Report, 2014). Thus, finding cleaner alternative energy sources is crucial. Of the renewable energy sources, solar power may be the most practical for meeting the global energy demand. It is estimated that enough energy from the sun hits the earth in an hour to meet the world’s energy use for a year (Katz, 2015). 2

For vehicles, alone, the world uses ~22,596,500 barrels of gasoline per day (U.S. Energy Information Administration, 2012). The problems with this usage include pollution (esp. carbon monoxide, non- methane hydrocarbon, oxides of nitrogen, volatile organic compounds, and heavy metals), 11,830.5 million metric tons of CO2 emissions per year (U.S. Energy Information Administration, 2012), finite and dwindling fossil fuel resources, rising costs, and dependence on imported oil. The impacts of transportation will only increase with population growth. The 1 “[] has led to atmospheric concentrations of carbon dioxide, methane and nitrous oxide that are unprecedented in at least the last 800,000 years.” – IPCC Fifth Assessment Synthesis Report, 2014

2 This theoretical amount is 350,000TW; the practical estimation is closer to 600TW (though estimates vary by 100-2500TW). Since current conversion efficiency is 10%, the estimates for onshore electricity production potential is ~60TW (Katz, 2015).


world’s 600 million light-duty vehicles may very well double by 2020 (MacLean, H. L., & Lave, L. B., 2003). 95% of these vehicles are dependent on oil (Mayyas et al., 2012), but with new technologies, such as hydrogen, this number will be able to decrease.

There are, at present, three main ways of harvesting solar energy: solar thermal, photovoltaic, and fuels. There are tradeoffs for each approach; price, efficiency, storage, and distribution are all major challenges for the adoption of solar power. Fuels may offer a cheaper alternative to the more efficient photovoltaic (Katz, 2015).

The process of hydrolysis, or ‘water splitting’, can make solar energy storable and transportable, in the form of Hydrogen (H) fuel (see Equation 1 and Figure 1), produces no CO2 (or other pollutants) in the production of H or the burning of H (Katz, 2015; Abu-Hamed, Karni, & Epstein, 2007), and therefore has the potential to make renewable energy more easily adopted as a substitute for fossil fuels (Katz, 2015; Abu-Hamed, Karni, & Epstein, 2007).

Metal + H2O Metal Oxide + H2 (1)

Figure 1: Hydrolysis reaction model

The principle chemical reaction in this experiment is:

2 Al + 3 H2O Al2O3 + 3 H2 (1)

(with NaOH as a catalyst)

According to recent studies, metals such as boron, zinc and aluminum can be used in hydrolysis of water. These studies also suggest that solar energy could be used to reverse the oxidation of metal, a by-product of the hydrolysis (Irina Vishnevetsky, 2008), recycling the metal for reuse in hydrogen production. Furthermore, aluminum in the process can come from waste aluminum, such as aluminum cans (Martınez et al., 2005).


However, as of 2005, no hydrogen storage technologies could compete with conventional internal combustion engines (ICE) in terms of volume constraints (and by extension vehicle ranges), and cost (Burke & Gardiner, 2005). Hydrogen has also been viewed critically because of its explosive nature, but on-board hydrogen production in vehicles could substitute hydrogen with a benign metal, making storage and transportation safer and reduce volume constraints (Abu-Hamed et al., 2007).

The aim of this research is to better understand the hydrolysis kinetics (speed) and efficiency of hydrogen production, as well as to investigate possible improvements in on-board hydrogen fuel production in vehicles. Specifically, this study will look at the effects of the catalyst, Sodium Hydroxide (NaOH), concentration on the rate of production of small amounts of hydrogen gas, through metal hydrolysis (see Equation 1). Finding the NaOH concentration that produces the most efficient reaction could solve some of the problems on-board H production faces (such as volume and weight constraints).

This type of research is important to the global transition from fossil fuels to solar power. Better understanding of the hydrolysis kinetics (speed) and best practices for efficient Hydrogen production, could increase the development on-board Hydrogen fuel production in vehicles.

Previous studies suggest that the catalyst concentration affects rates of hydrogen production in metal hydrolysis (Stern, 2014). The study should determine how rates of hydrogen production compare among NaOH solution concentrations. It is expected that the rate of hydrogen production in metal hydrolysis will be affected by NaOH solution concentrations and that higher concentrations of NaOH solution will increase the rate of production, as seen in similar trials (Stern, 2014).

Methods:

The study was conducted in the Arava Institute for Environmental Studies Laboratory, at Kibbutz Ketura, Israel. Throughout all times in the study, the temperature in the lab was kept at 25C. Through a series of repetitions of the metal hydrolysis process, the tests measured hydrogen production rates at different NaOH solution concentrations of 0.1M, 0.5M, 1M, and 5M.

Chemical Reactants

Related studies have found Aluminum, as well as Boron, to be promising candidates for hydrogen production through metal hydrolysis (Abu-Hamed, 2007). A similar study has also suggested the superiority of NaOH as a catalyst in such a reaction (Stern, 2014).

The equation below used to calculate the mass of NaOH used to produce different NaOH concentrations:


M= nV

=m /MWV

(2)

Where M is concentration (mol/L), n is number of moles (mol), V is volume (L), m is the mass (g), and MW is molecular weight of NaOH (40 g/mol). Calculations of grams of NaOH per 100mL of water needed for each concentration are as follows:

0.1M= nV

= m /400.1 liter

m= 0.4 g NaOH

0.5M= nV

= m/ 400.1liter

m= 2.0 g NaOH

1.0M= nV

= m /400.1 liter

m= 4.0g NaOH

2.0M= nV

= m / 400.1liter

m=8.0g NaOH

The mass of Aluminum needed for the chemical reaction was calculated with Equation 1.

2 Al + 3 H2O Al2O3 + 3 H2 (1)

(with NaOH as a catalyst)

According to Equation 1, 2 moles of Aluminum produce 3 moles of Hydrogen. With the assumption of the reaction produces 50 mL of Hydrogen gas, the moles of Hydrogen were calculated with the following equation:

n=V/MW (3)

Where n is the number of Hydrogen moles, V is the Hydrogen volume in liters, and the molar mass of Hydrogen is 22.4 g/mol.

n=0.050/22.4 =0.0022 moles of H

Since 2 moles of Aluminum produce 3 moles of Hydrogen, as seen in Equation 1, the number of aluminum moles were calculated by multiplying the number of moles of Hydrogen by 2/3. This resulted in 0.0015 moles of Al, which were converted into mass using Equation 3.

m=0.0015*27=0.04 grams of Al

Apparatus


A calibrated mass scale was used to measure 0.4g of NaOH and 0.04 grams of Al. NaOH was dissolved into a flask filled with 100mL of distilled water. Care was taken when spooning out contents onto wax paper to reduce time in which NaOH was exposure to air. Al was in the form of aluminum foil. Aluminum foil was cut into 4cm x 2cm rectangles folded three times.

The reactions will take place in an airtight glass jar attached to flexible rubber tubing. The tubing will be connected to a buret, the bottom end of which will be submerged in water. Water was extracted to fill a buret tube to the 50mL line. As the reaction takes place, the hydrogen gas produced will displace water in the buret tube. Airtight glass jars were filled with enough NaOH solution to submerge the aluminum foil.

Fig. 2: Testing apparatus; (left to right) Mass scale, NaOH (with red cap), notebook (for recording water levels and qualitative notes), rubber gloves (for handling

NaOH, Al foil, and byproducts), sitting atop each jack is an airtight glass jar with rubber tubing running through the water in the beakers to the submerged bottom

end of the burets (which are clamped to upright poles behind)

Data collection

Two identical apparatuses were set up to obtain two simultaneous and similar repetitions. This process was repeated for NaOH solutions with a concentration of 0.5M, 1M and 5M.


Rate of Hydrogen production was recorded at intervals of 5 minutes for NaOH solutions with a concentration of 0.1M; 2 minute intervals for 0.5M, 1M, and 5M because of the increased rate of H production. Timekeeping was started when aluminum foil was added to the solution in the first airtight glass jar and the lid closed. Airtight glass jars were swirled ever 5 minutes or so, to encourage aluminum foil to completely submerge. Water levels were observed by eye and recorded by hand. Quantitative recording ended when water levels reached 0mL, after which some qualitative observations were made.

Data Analysis

The effects of the variable concentration on rates of production were compared. Results were compared with data from the Stern study of 2014, as well as other studies of on- board H production technologies and external H production technologies. Lastly, the practicality of application of this hydrogen production method for on- onboard vehicles was assessed with comparisons to the capabilities of gasoline cars.

Results

The results show a positive correlation between NaOH concentration and H production rate (Fig. 1). The most rapid rate was seen in the NaOH concentration of 2.0M, as expected.

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0.1M0.5M1.0M2.0M

Time (minutes)

Ave

rage

Vol

um

e of

H P

rod

uce

d

(mL)

Fig. 3: Hydrogen Production Rate for various NaOH concentrations


This study found the highest rate of H production to be 5.4 mL/min for the 2.0M NaOH concentrations (Fig. 4 and Table 1). The next fastest rate of H production was 1.0M at 1.3 mL /min. The rate for 0.5M was 0.9 mL/min and for 0.1M was 0.6 mL/min.

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f(x) = 0.648214285714286 x − 1.79821428571429

f(x) = 0.901468985507247 x − 3.97846641821947f(x) = 1.34538069806461 x − 4.5464376019094f(x) = 5.42642857142857 x − 9.96690476190476

0.1MLinear (0.1M)0.5MLinear (0.5M)1.0MLinear (1.0M)2.0MLinear (2.0M)

Time (minutes)

Ave

rage

Vol

um

e of

H P

rod

uce

d (

mL)

Fig. 4: Average Hydrogen Production Rate for various NaOH concentrations

0.1M 0.5M 1.0M 2.0M

Avg. Ratesy=0.6482x -

1.7982y=0.9015x -

3.9785y=1.3454x -

4.5464y=5.4264x -

9.9669Calculated

finish times (mins) 79.91082999 59.87631725 40.54288687 11.05095459


Table 1: Comparison of Average rates and Calculated finish times

Discussion

Results from this experiment, alone, are not enough – more repetitions would be needed to confirm the trends, however, in testing, some trends became evident and these trends are supported by results found in a similar study (Stern, 2014). NaOH concentration did affect the rate of H production in metal hydrolysis and positively correlated (Fig. 3), as predicted. It is this positive correlation that was seen previously by Stern (2014) (Appendix).

The unexpected element in the results was the difference between rates. The rate for 2.0M, the fastest rate, was about 400% faster than the rate for 1.0M, which is half of the concentration. The rate for 2.0M is also around 600% of the rate for 0.5M, which is only a 400% increase in concentration.

This significant increase in rate of hydrogen production may have resulted from improperly cleaned equipment in repetitions of 0.1M, 0.5M, and half of the repetitions for 1.0M. After testing 0.1M, 0.5M, and half of the repetitions of 1.0M, an orange substance found inside Jar 1 lid (Fig. 5). After cleaning lids and jars, H production times for 1.0M sped up. Orange substance is likely Al2O3, a byproduct of the reaction. This substance may have blocked H gas flow or it may have inhibited the reaction by decreasing unreacted aluminum surface area, as seen in studies on boron (Wahbeh et al., 2012).


Fig. 5: Both test jars after third trial (May 5, 2015). Byproducts of reaction - blackened liquid and corroded aluminum foil are visible, as well as orange

substance found inside Jar 1 lid.

Another trend that became evident was the slower rates seen in Jar 1 in all tests but those preformed on 0.1M NaOH concentrations. Because the discrepancy was still seen in testing after the lids and jars were cleaned, though to a much lesser extent, and because the difference in jar rates was not observed in testing of 0.1M, it is unlikely that the orange substance was the only cause. There is a possibility that the differences come from recording errors; visual observation of water levels was done with two sets of eyes when testing 0.1M, and one set for other concentrations, with Jar 1 being observed first consistently (though this may have been counteracted by roughly the same true start time lapses, as aluminum was added to Jar 1 and then Jar 2). There may also be the possibility of the appearance of a small leakage in Jar 1, though more testing would be needed to confirm this.

If this research were to continue, testing to attempt to understand these two deviations may lead to a better understanding of the process, however testing molarities above 2.0M and at what concentration effects begin to wane, may be more pertinent to the objectives of this study. This would need to be done with either a smaller amount of aluminum or a more exact way to record levels, as the rates are expected to increase dramatically.

Still, even with molarities above 2.0M, it is not predicted that there will be a sufficient increase in rate to outperform other methods of hydrogen production. This study found the highest rate of H production to be 5.4 mL/min for the 2.0M NaOH concentrations and the next fastest rate to be 1.3 mL /min for 1.0M (Fig. 4 and Table 1). Stern found rates of ~4.9 mL/min for a 2.0M concentration of NaOH combined with 0.032g of Al (as compared with 0.04g of Al in this study) at 25C (2014). While the rates found in this study for 2.0M are slightly faster than those found by Stern, the rates Stern saw at 50C were much faster than either experiments at 25C. At 50C, rates were seen to be ~16.6 mL/min for 2.0M and ~15.5 mL/min for 1.0M (Stern, 2014). The trend of increased rates of H production with higher temperatures has been seen with other metals and processes (Gálvez et al., 2008; Stern, 2014; Vishnevetsky et al., 2011; Wahbeh et al., 2012; Weidenkaff et al., 2000). The rates found in this study are also far below those found for boron at very high temperatures has been seen at ~70 ccm, at highest, 6.5 minutes into the reaction (Wahbeh et al., 2012).

Considering that ~56,000,000 mL of H (5kg) are needed to drive a car 500km (Abu-Hamed et al., 2007), comparable with gasoline- powered vehicles, over a million times more H would need to be produced. Though it’s unclear how this type of expansion would affect rates of production, it is evident such an augmentation of materials involved in the reaction would make the technology nonviable in application, on-board vehicles.

Variables other than concentration, such as higher temperatures, metal type, reactant quantity, and ways to increase surface area, might produce greater increases in hydrogen production rate and should be studied.


Conclusion

Hydrolysis, or ‘water splitting’, emits no pollution, and recent studies suggest that solar energy could be used to recycle the metal oxidized in the process, which could close the loop of this process and make a storable and transportable, in the form of Hydrogen (H) fuel form of solar energy, that is more competitive with fossil fuels. This study investigated efficient hydrogen production, for use on-board vehicles, by measuring the effects of the catalyst, Sodium Hydroxide (NaOH), concentration on the rate of production of hydrogen gas. The time required to produce 50mL of H gas was recorded for exactly calculated amounts of Water, Aluminum foil, and NaOH. A positive correlation was seen between NaOH concentration and H production rate, and fastest rate was found in the NaOH concentration of 2.0M (significantly higher than other tested concentrations). Rates found in this study were not as high as those found with higher temperatures (Stern, 2014) and with different metals and temperatures (Wahbeh et al., 2012). While this method may not lead be a practical alternative to fossil fuels, there are other promising on-board hydrogen production methods may produce greater increases than seen in this study in hydrogen production rate and should be studied.

Acknowledgements: Tareq Abu Hamed, Gabi, Avi, Jess, Dan, Alex, Arava Institute for Environmental Sciences, Kibbutz Ketura

References:

Abu-Hamed, T., Karni, J., & Epstein, M. (2007). The use of boron for thermochemical storage and distribution of solar energy. Solar Energy, 81:93-101.

Wahbeh,B., Abu Hamed,T., & Kasher, R. (2012). Hydrogen and boric acid production via boron hydrolysis, Renewable Energy, 48:10-15.

Burke, A., & Gardiner, M. (2005). “Hydrogen Storage Options: Technologies and Comparisons for Light-duty Vehicle Applications.” Hydrogen Pathways Program Institute of Transportation Studies University of California-Davis, UCD-ITS-RR-05-01.

Gálvez, M.E., Frei, A., Albisetti, G., Lunardi, G., & Steinfeld, A. (2008). Solar hydrogen production via a two-step thermochemical process based on MgO/Mg redox reactions—thermodynamic and kinetic analyses. Int J Hydrogen Energy, 33: 2880–2890.


Katz, J. (2015). “Global Energy: Where in World Will it Come From?” [Seminar lecture and PowerPoint]. Denison University.

MacKay, D.J.C. (2009). Sustainable Energy - without the hot air. Cambridge, England: UIT.

MacLean, H. L., & Lave, L. B. (2003). Life cycle assessment of Automobile/Fuel options. Environmental Science & Technology, 37: 5445-5452.

Martınez, S.S., Benites, W.L., Gallegos, A.A., & Sebastian, P.J. (2005). Recycling of aluminum to produce green energy. Solar Energy Materials and Solar Cells, 88: 237-243.

Mayyas, A., Qattawi, A., Omar, M., & Shan, D. (2012). Design for sustainability in automotive industry: A comprehensive review. Renewable and Sustainable Energy Reviews, 16(4), 1845-1862.

Stern, T. “Hydrogen Production by Low Temperature Metal Hydrolysis.” Arava Institute for Environmental Studies. Advisor: Dr. Tareq Abu Hamed. Ketura, 2014.

Vishnevetsky, I., Epstein, M., Abu-Hamed, T., & Karni, J. (2008). Boron hydrolysis at moderate temperatures: First step to solar fuel cycle for transportation. Journal of Solar Energy Engineering (Transactions of the ASME), 130(1)

Vishnevetsky, I., Berman, A., & Epstein, M. (2011). Features of solar thermochemical redox cycles for hydrogen production from water as a function of reactants’ main characteristics. Int J Hydrogen Energy, 36: 2817–2830

Weidenkaff, A., Reller, A., Wokaun, A., & Steinfeld, A. (2000). Thermogravimetric analysis of the ZnO/Zn water splitting cycle. Thermochimical Acta, 359: 69–75

Yavor, Y., Goroshin, S., Bergthorson, J.M., Frost, D.L., Stowe, R., and Ringuette, S. (2013). Enhanced hydrogen generation from aluminum water reactions. International Journal of Hydrogen Energy 38: 14992–15002.


Appendices

(Stern, 2014)

This

study:

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