Powering our Cars: An Analysis of Various Fuel Sources in Environmental and Green Chemistry
- Page ID
- 418936
\( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)
\( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)
\( \newcommand{\dsum}{\displaystyle\sum\limits} \)
\( \newcommand{\dint}{\displaystyle\int\limits} \)
\( \newcommand{\dlim}{\displaystyle\lim\limits} \)
\( \newcommand{\id}{\mathrm{id}}\) \( \newcommand{\Span}{\mathrm{span}}\)
( \newcommand{\kernel}{\mathrm{null}\,}\) \( \newcommand{\range}{\mathrm{range}\,}\)
\( \newcommand{\RealPart}{\mathrm{Re}}\) \( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)
\( \newcommand{\Argument}{\mathrm{Arg}}\) \( \newcommand{\norm}[1]{\| #1 \|}\)
\( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)
\( \newcommand{\Span}{\mathrm{span}}\)
\( \newcommand{\id}{\mathrm{id}}\)
\( \newcommand{\Span}{\mathrm{span}}\)
\( \newcommand{\kernel}{\mathrm{null}\,}\)
\( \newcommand{\range}{\mathrm{range}\,}\)
\( \newcommand{\RealPart}{\mathrm{Re}}\)
\( \newcommand{\ImaginaryPart}{\mathrm{Im}}\)
\( \newcommand{\Argument}{\mathrm{Arg}}\)
\( \newcommand{\norm}[1]{\| #1 \|}\)
\( \newcommand{\inner}[2]{\langle #1, #2 \rangle}\)
\( \newcommand{\Span}{\mathrm{span}}\) \( \newcommand{\AA}{\unicode[.8,0]{x212B}}\)
\( \newcommand{\vectorA}[1]{\vec{#1}} % arrow\)
\( \newcommand{\vectorAt}[1]{\vec{\text{#1}}} % arrow\)
\( \newcommand{\vectorB}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)
\( \newcommand{\vectorC}[1]{\textbf{#1}} \)
\( \newcommand{\vectorD}[1]{\overrightarrow{#1}} \)
\( \newcommand{\vectorDt}[1]{\overrightarrow{\text{#1}}} \)
\( \newcommand{\vectE}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash{\mathbf {#1}}}} \)
\( \newcommand{\vecs}[1]{\overset { \scriptstyle \rightharpoonup} {\mathbf{#1}} } \)
\(\newcommand{\longvect}{\overrightarrow}\)
\( \newcommand{\vecd}[1]{\overset{-\!-\!\rightharpoonup}{\vphantom{a}\smash {#1}}} \)
\(\newcommand{\avec}{\mathbf a}\) \(\newcommand{\bvec}{\mathbf b}\) \(\newcommand{\cvec}{\mathbf c}\) \(\newcommand{\dvec}{\mathbf d}\) \(\newcommand{\dtil}{\widetilde{\mathbf d}}\) \(\newcommand{\evec}{\mathbf e}\) \(\newcommand{\fvec}{\mathbf f}\) \(\newcommand{\nvec}{\mathbf n}\) \(\newcommand{\pvec}{\mathbf p}\) \(\newcommand{\qvec}{\mathbf q}\) \(\newcommand{\svec}{\mathbf s}\) \(\newcommand{\tvec}{\mathbf t}\) \(\newcommand{\uvec}{\mathbf u}\) \(\newcommand{\vvec}{\mathbf v}\) \(\newcommand{\wvec}{\mathbf w}\) \(\newcommand{\xvec}{\mathbf x}\) \(\newcommand{\yvec}{\mathbf y}\) \(\newcommand{\zvec}{\mathbf z}\) \(\newcommand{\rvec}{\mathbf r}\) \(\newcommand{\mvec}{\mathbf m}\) \(\newcommand{\zerovec}{\mathbf 0}\) \(\newcommand{\onevec}{\mathbf 1}\) \(\newcommand{\real}{\mathbb R}\) \(\newcommand{\twovec}[2]{\left[\begin{array}{r}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\ctwovec}[2]{\left[\begin{array}{c}#1 \\ #2 \end{array}\right]}\) \(\newcommand{\threevec}[3]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\cthreevec}[3]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \end{array}\right]}\) \(\newcommand{\fourvec}[4]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\cfourvec}[4]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \end{array}\right]}\) \(\newcommand{\fivevec}[5]{\left[\begin{array}{r}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\cfivevec}[5]{\left[\begin{array}{c}#1 \\ #2 \\ #3 \\ #4 \\ #5 \\ \end{array}\right]}\) \(\newcommand{\mattwo}[4]{\left[\begin{array}{rr}#1 \amp #2 \\ #3 \amp #4 \\ \end{array}\right]}\) \(\newcommand{\laspan}[1]{\text{Span}\{#1\}}\) \(\newcommand{\bcal}{\cal B}\) \(\newcommand{\ccal}{\cal C}\) \(\newcommand{\scal}{\cal S}\) \(\newcommand{\wcal}{\cal W}\) \(\newcommand{\ecal}{\cal E}\) \(\newcommand{\coords}[2]{\left\{#1\right\}_{#2}}\) \(\newcommand{\gray}[1]{\color{gray}{#1}}\) \(\newcommand{\lgray}[1]{\color{lightgray}{#1}}\) \(\newcommand{\rank}{\operatorname{rank}}\) \(\newcommand{\row}{\text{Row}}\) \(\newcommand{\col}{\text{Col}}\) \(\renewcommand{\row}{\text{Row}}\) \(\newcommand{\nul}{\text{Nul}}\) \(\newcommand{\var}{\text{Var}}\) \(\newcommand{\corr}{\text{corr}}\) \(\newcommand{\len}[1]{\left|#1\right|}\) \(\newcommand{\bbar}{\overline{\bvec}}\) \(\newcommand{\bhat}{\widehat{\bvec}}\) \(\newcommand{\bperp}{\bvec^\perp}\) \(\newcommand{\xhat}{\widehat{\xvec}}\) \(\newcommand{\vhat}{\widehat{\vvec}}\) \(\newcommand{\uhat}{\widehat{\uvec}}\) \(\newcommand{\what}{\widehat{\wvec}}\) \(\newcommand{\Sighat}{\widehat{\Sigma}}\) \(\newcommand{\lt}{<}\) \(\newcommand{\gt}{>}\) \(\newcommand{\amp}{&}\) \(\definecolor{fillinmathshade}{gray}{0.9}\)This Exemplar will teach the following concepts from the ACS Examinations Institute General Chemistry ACCM:
- II. D. 1. b. Bond dissociation enthalpies can be used to estimate the change in enthalpy for a reaction.
- VI. C. 1. a. Heat flow into the system is defined as endothermic; heat flow out of the system is defined as exothermic.
- VI. C. 2. a. Enthalpies of reaction, such as those incorporated in a thermochemical equation, can be used to calculate energy flow for a given amount of reactant or product.
- IX. E. 2. b. Many hydrocarbons can be used as fuels; the combustion of a hydrocarbon in air produces carbon dioxide and water.
The Thermodynamics of Fuel
Enthalpy (H) is defined as the sum of the internal energy of a thermodynamic system and the product of pressure and volume. It describes the current state of the system and is measured in units of energy (J). When the pressure and volume of a system are held constant, the change in enthalpy (ΔH) represents the movement of heat in or out of the system.1 For example, the ΔH for ice melting into liquid water describes the amount of heat added to the ice in order for it to change state, and ΔH will be positive because liquid water has greater enthalpy than ice, and ΔH = Hproducts - Hreactants. Furthermore, ΔH is directly proportional the amount of the substance changing in the system, so it has units J/mol. Chemical reactions all have a ΔH value, which describes whether the reaction releases energy and is exothermic (ΔH < 0) or requires energy and is endothermic (ΔH > 0). Chemical reactions will release energy into the environment when the bonds in the product molecules are stronger than bonds in the reactant molecules.
The substances that power our vehicles are chosen to be fuels because of their property to react exothermically in a chemical reaction. In effect, energy is stored in their chemical bonds, and can be released in a controlled manner under the correct conditions and activation energy. Better, more energy-dense fuels will be able to release more energy for each unit consumed. The fuels discussed in this project are consumed through combustion, a type of chemical reaction in which a substance containing carbon, hydrogen, and oxygen atoms, reacts with oxygen under pressure or heat to create water, carbon dioxide (CO2), and a large amount of heat.2 Combustion reactions are always exothermic because the bond dissociation energies in the products, especially the double bonds between carbon and oxygen in CO2, are much higher than the bond dissociation energies in the hydrocarbon or elemental oxygen.2 Different fuels for combustion engines in cars are chosen not only based on their energy density, but also based on the cost of production, and ideal reaction conditions.
|
Gasoline/E10 |
Diesel |
E85 |
CNG |
Hydrogen |
|
|---|---|---|---|---|---|
|
Chemical Structure |
Hydrocarbons, C4 through C12, with 10% EtOH |
Hydrocarbons, C8 through C25 |
EtOH + Hydrocarbons of Gasoline |
Majority CH4, but also C2H6 |
H2 |
|
Physical State |
Liquid |
Liquid |
Liquid |
Gas |
Gas or Liquid |
|
Energy Content (lower heating value) |
112,144-116,090 Btu/gal |
128,488 Btu/gal |
83,950-95,450 Btu/gal |
21,160 Btu/lb |
51,585 Btu/lb |
|
Maintenance Issues |
N/A |
N/A |
Similar to conventional fuels, however special lubricants for the engine may be required |
High-pressure tanks are required |
High-pressure tanks are required, although this problem is eliminated with the use of fuel cells |
|
Cost ($/Million Btu*) |
$41.12 |
$43.82 |
$58.22 |
$24.15 |
Data not available |
*Btu = British Thermal Units: Heat required to raise temperature of 1 pound of water by 1 degree Fahrenheit
Energy Value of Gasoline, Ethanol, and Natural Gas
Gasoline
Gasoline is derived from petroleum, a yellow-black liquid fossil fuel that is mainly composed of hydrocarbons, and is by far the most common fuel for cars around the world. To enhance the stability, efficiency, and performance of gasoline, various blending agents and additives are added to the gasoline hydrocarbon mixture to make for an average of 150 distinct compounds in finished gasoline. Such compounds include anti-rust agents, anti-oxidants, upper-cylinder lubricants, anti-icing agents, among many more.5 However, for the sake of simplicity, the energy value of octane (C8H18) will be used in the calculation, which is a good representation of the energy properties of gasoline.
Figure 1: An industrial process plant in Washington where petroleum is refined6
Figure 2: Ball-and-stick representation of 2,2,4-Trimethylpentane, or Isooctane7
The combustion of octane can be represented in the following balanced equation:
C8H18 + 25/2 O2 → 8CO2 + 9H2O
The standard enthalpy change of formation (ΔHf) for a chemical reaction occurring at standard temperature and pressure can be calculated by the following equation:
ΔH°rxn = ΣnΔH°f (products) - ΣnΔH°f (reactants)
Given that the ΔHf of octane, carbon dioxide, water, and oxygen are -250 kJ/mol, -394 kJ/mol, -286 kJ/mol, and 0 kJ/mol respectively, the change in enthalpy (ΔH) can be calculated by:
ΔHf = [(8 mol CO2 x -394 kJ/mol) + (9 mol H2O x -286 kJ/mol)] - [(1 mol C8H18 x -250 kJ/mol) + (25/2 mol O2 x 0 kJ/mol)] = -5726 kJ/mol - -250 kJ/mol = -5476 kJ/mol
Using this value, as well as the density and molar mass of octane, the heat released by combusting one gallon of octane can be calculated:
C = 8 x 12.011 = 96.088
H = 18 x 1.008 = 18.144
Molar Mass = 96.088 + 18.144 = 114.23 grams/mol
Density of Octane (C8H18) = 0.737 g/mL
q = 1 gallon C8H18 * (3.79L/1 gallon C8H18) * (1000 mL/1 L) * (0.737 g/1 mL) * (1 mol C8H18/114.23 g) * (-5476 kJ/1 mol C8H18) = -133,903 kJ
Ethanol
Ethanol (C2H6O), an organic compound derived from crop waste, sugarcane, corn, or wood chips, is increasingly being added to gasoline today to enhance its octane rating. Gasoline with a higher octane rating is more stable and withstands detonation to a greater extent, allowing for it to burn rather than explode.8 However, increasing a gasoline’s octane level is an expensive process, which is why gasoline with higher octane rating is considered “premium” fuel and is more costly. Ethanol stabilizes low-grade gasoline in a more financially feasible manner, allowing for its rapid widespread use throughout the United States. Currently, more than 98% of United States Gasoline utilizes E10, a fuel mixture containing 10% ethanol and 90% gasoline.9
Figure 3: Warning label for higher than expected ethanol content in gasoline10
The combustion of ethanol can be represented in the following balanced equation:
C2H6O(l) + 3O2(g) → 2CO2(g) + 3H2O(l)
Given that the ΔHf of ethanol, carbon dioxide, and water are -278 kJ/mol, -394 kJ/mol, and -286 kJ/mol respectively, the change in enthalpy (ΔH) can be calculated by:
ΔHf = [(2 mol CO2 x -394 kJ/mol) + (3 mol H2O x -286 kJ/mol)] - [(1 mol C2H6O x -278 kJ/mol) + (3 mol O2 x 0 kJ/mol)] = -1646 kJ/mol - -278 kJ/mol = -1368 kJ/mol
Using this value, as well as the density and molar mass of ethanol, the heat released by combusting one gallon of ethanol can be calculated:
C = 2 x 12.011 = 24.022
H = 6 x 1.008 = 6.048
O = 1 x 15.999 = 15.999
Molar Mass = 24.022 + 6.048 + 15.999 = 46.039 grams/mol
Density of Ethanol (C2H6O) = 0.789 g/mL
q = 1 gallon C2H6O * (3.79L/1 gallon C2H6O) * (1000 mL/1 L) * (0.789 g/1 mL) * (1 mol C2H6O/46.07 g) * (-1368 kJ/1 mol C2H6O) = -88,794 kJ
By using the ratios of the fuel composition, the sum of the heat released from ethanol can be multiplied by 0.1 and the heat released from gasoline can be multiplied by 0.9 to yield a result of -129410 kJ.
qoctane = -133,903 kJ
qethanol = -88,794 kJ
qe10 = 0.1(qethanol) + 0.9(qoctane) = 0.1(-88,794 kJ) + 0.9(-133,904 kJ) = -129,410 kJ
Compared to the previously calculated released heat from gasoline, it can be observed E10 fuel releases less heat by 3.36%. This echoes a common concern by many Americans11; E10 fuel is less efficient and consequently results in lower mileage (making for a lesser dollar/mile ratio), not to mention, fosters damage to the fuel systems and the environment11.
100 - |(-129,410 kJ/-133,903 kJ) * 100| = 3.36%
Natural Gas
Although not as common, natural gas vehicles are growing rapidly in use both in the United States and other predominantly first world countries, with over 23 million vehicles worldwide.12 Two types of natural gas vehicles are manufactured depending on its driving range needs, that is, compressed natural gas (CNG) vehicles and liquefied natural gas (LNG) vehicles. CNG vehicles maintain natural gas in a gaseous state under pressure in a tank, while LNG vehicles maintain natural gas in a liquid form in a relatively smaller tank, allowing for it to yield a greater energy density than that of CNG vehicles. Although natural gas is composed of a number of hydrocarbons amongst other additives, it is composed of approximately 85-90% methane (CH4)3, which will be utilized in our calculation as a representation of this fuel. Specifically, the density of liquid methane will be used, meaning the calculation will display the heat released by an LNG vehicle.
Figure 4: Gas tanks in the trunk of a CNG vehicle13
The combustion of methane can be represented in the following balanced equation:
CH4(g) + 2O2(g) → CO2(g) + 3H2O(g)
Given that the ΔHf of methane, carbon dioxide, and water are -75 kJ/mol, -394 kJ/mol, and -286 kJ/mol respectively, the change in enthalpy (ΔH) can be calculated by:
ΔHf = [(1 mol CO2 x -394 kJ/mol) + (2 mol H2O x -286 kJ/mol)] - [(1 mol CH4 x -75 kJ/mol) + (2 mol O2 x 0 kJ/mol)] = -966 kJ/mol - -75 kJ/mol = -891 kJ
Using this value, as well as the liquid density and molar mass of methane, the heat released by combusting one gallon of methane can be calculated:
C = 1 x 12.011 = 12.011
H = 4 x 1.008 = 4.032
Molar Mass = 12.011 + 4.032 = 16.043 grams/mol
Density of Ethanol (CH4) = 0.4228 g/mL
q = 1 gallon CH4 * (3.79L/1 gallon CH4) * (1000 mL/1 L) * (0.4228 g/1 mL) * (1 mol CH4/16.043 g) * (-891 kJ/1 mol C2H6O) = -88,995 kJ
Economy and Efficiency
With one gallon of natural gas releasing -88,995 kJ (see Table 1), it proves to be a less efficient source of fuel than that of gasoline. Although the combustion of pure ethanol releases -88,794 kJ, a mere difference in 201 kJ, it is important to note that the densest fuel mixture of ethanol for cars in the USA includes only 85% ethanol and 15% gasoline, a fuel mixture known as E85. However, as seen in the E10 mixture, which yielded 3.36% less heat than that of pure gasoline, increasing the ethanol content sacrifices fuel economy.
As of July 2022, the average gasoline cost in the United States is $4.70 per gallon, and the average LNG cost was $3.15.4 Since the most commonly used fuel in the United States is E10, by using the heating value for E10 (-129410.1 kJ), it can be found that E10 yields a value of 27534.06 kJ/$ while LNG yields a value of 28252.38 kJ/$. In this manner, natural gas displays a greater fuel economy (with the current market prices), but with the lack of fueling stations, the relatively mediocre fuel efficiency, and the greater consumer cost of CNG/LNG cars, they are not as widespread12.
References
1. The Editors of Encyclopaedia Britannica. Enthalpy. https://www.britannica.com/science/enthalpy (accessed Nov 10, 2022).
2. Schmidt-Rohr, K. Why Combustions Are Always Exothermic, Yielding about 418 KJ per Mole of O2. Journal of Chemical Education 2015, 92 (12), 2094–2099. 3. Vehicle Technologies Office, U.S. Department of Energy. Fuel properties comparison. https://afdc.energy.gov/fuels/properties (accessed Nov 10, 2022).
4. Bourbon, E. Clean Cities Alternative Fuel Price Report, July 2022. https://afdc.energy.gov/publications/ (accessed Nov 10, 2022). 5. Ritter, S. Chemical & Engineering News. American Chemical Society February 21, 2005.
5. Ritter, S. Chemical & Engineering News. American Chemical Society February 21, 2005.
6. Wikimedia Commons contributors. File:Anacortes Refinery 31911.JPG - Wikimedia Commons. https://commons.wikimedia.org/wiki/F...nery_31911.JPG (accessed Dec 8, 2022).
7. National Center for Biotechnology Information. https://pubchem.ncbi.nlm.nih.gov/com...n=3D-Conformer (accessed Dec 7, 2022).
8. Stolark, J. Fact sheet: A brief history of octane in gasoline: From lead to ethanol. https://www.eesi.org/papers/view/fac...tory-of-octane (accessed Nov 10, 2022).
9. Oak Ridge National Laboratory. Ethanol. https://www.fueleconomy.gov/feg/ethanol.shtml (accessed Nov 11, 2022).
10. Wikimedia Commons contributors. File:EPA E15 warning label.jpg - Wikimedia Commons. https://commons.wikimedia.org/wiki/F...ning_label.jpg (accessed Dec 8, 2022).
11. Performance, B. The disadvantages of adding ethanol to your fuel. https://www.bellperformance.com/blog...l-to-your-fuel (accessed Dec 7, 2022).
12. Vehicle Technologies Office, U.S. Department of Energy. Natural gas vehicles. https://afdc.energy.gov/vehicles/natural_gas.html (accessed Dec 7, 2022).
13. Wikimedia Commons contributors. File:SAO 09 2008 Fiat Siena tetrafuel 2 views v1.jpg. https://commons.wikimedia.org/wiki/F...2_views_v1.jpg (accessed Dec 8, 2022).
All thermodynamic data was retrieved from the NIST Chemistry WebBook: https://webbook.nist.gov/chemistry/