To get the enthalpy of combustion for 1 mole of acetylene, divide the balanced equation by 2 C2H 2(g) + 5 2 O2(g) 2CO2(g) + H 2O(g) Now the expression for the enthalpy of combustion will be H comb = (2 H 0 CO2 +H H2O) (H C2H2) H comb = [2 ( 393.5) +( 241.6)] (226.7) H comb = 1255.3 kJ Because enthalpy is a state function, a process that involves a complete cycle where chemicals undergo reactions and are then reformed back into themselves, must have no change in enthalpy, meaning the endothermic steps must balance the exothermic steps. After that, add the enthalpies of formation of the products. The value of a state function depends only on the state that a system is in, and not on how that state is reached. And we're multiplying this by five. This is one version of the first law of thermodynamics, and it shows that the internal energy of a system changes through heat flow into or out of the system (positive q is heat flow in; negative q is heat flow out) or work done on or by the system. It is the heat evolved when 1 mol of a substance burns completely in oxygen at standard conditions. Hess's Law states that if you can add two chemical equations and come up with a third equation, the enthalpy of reaction for the third equation is the sum of the first two. We still would have ended This H value indicates the amount of heat associated with the reaction involving the number of moles of reactants and products as shown in the chemical equation. By definition, the standard enthalpy of formation of an element in its most stable form is equal to zero under standard conditions, which is 1 atm for gases and 1 M for solutions. We will include a superscripted o in the enthalpy change symbol to designate standard state. Algae can produce biodiesel, biogasoline, ethanol, butanol, methane, and even jet fuel. (Note: You should find that the specific heat is close to that of two different metals. So to represent those two moles, I've drawn in here, two molecules of CO2. Here is a less straightforward example that illustrates the thought process involved in solving many Hesss law problems. The molar heat of combustion \(\left( He \right)\) is the heat released when one mole of a substance is completely burned. For more tips, including how to calculate the heat of combustion with an experiment, read on. Measure the mass of the candle and note it in g. When the temperature of the water reaches 40 degrees Centigrade, blow out the substance. Calculate the molar heat of combustion. Include your email address to get a message when this question is answered. To get kilojoules per mole These values are especially useful for computing or predicting enthalpy changes for chemical reactions that are impractical or dangerous to carry out, or for processes for which it is difficult to make measurements. To log in and use all the features of Khan Academy, please enable JavaScript in your browser. The Experimental heat of combustion is inaccurate because it does not factor in heat loss to surrounding environment. single bonds over here, and we show the formation of six oxygen-hydrogen consent of Rice University. (i) ClF(g)+F2(g)ClF3(g)H=?ClF(g)+F2(g)ClF3(g)H=? Explain why this is clearly an incorrect answer. Using enthalpies of formation from T1: Standard Thermodynamic Quantities calculate the heat released when 1.00 L of ethanol combustion. So for the combustion of one mole of ethanol, 1,255 kilojoules of energy are released. \[30.0gFe_{3}O_{4}\left(\frac{1molFe_{3}O_{4}}{231.54g}\right) \left(\frac{-3363kJ}{3molFe_{3}O_{4}}\right) = -145kJ\], Note, you could have used the 0.043 from step 2, Enthalpies of formation are usually found in a table from CRC Handbook of Chemistry and Physics. Hcomb (H2(g)) = -276kJ/mol, Note, in the following video we used Hess's Law to calculate the enthalpy for the balanced equation, with integer coefficients. One box is three times heavier than the other. , Calculate the grams of O2 required for the combustion of 25.9 g of ethylcyclopentane, A 32.0 L cylinder containing helium gas at a pressure of 38.5 atm is used to fill a weather balloon in order to lift equipment into the stratosphere. Creative Commons Attribution/Non-Commercial/Share-Alike. . Estimate the heat of combustion for one mole of acetylene: C2H2 (g) + O2 (g) 2CO2 (g) + H2O (g) Bond Bond Energy/ (kJ/mol CC 839 C-H 413 O=O 495 C=O 799 O-H 467 A. The distance you traveled to the top of Kilimanjaro, however, is not a state function. And that would be true for while above we got -136, noting these are correct to the first insignificant digit. tepwise Calculation of \(H^\circ_\ce{f}\). (a) What is the final temperature when the two become equal? We also can use Hesss law to determine the enthalpy change of any reaction if the corresponding enthalpies of formation of the reactants and products are available. then you must include on every digital page view the following attribution: Use the information below to generate a citation. Its unit in the international system is kilojoule per mole . Right now, we're summing Describe how you would prepare 2.00 L of each of the following solutions. Subtract the reactant sum from the product sum. 0.250 M NaOH from 1.00 M NaOH stock solution. 3 Put the substance at the base of the standing rod. However, we often find it more useful to divide one extensive property (H) by another (amount of substance), and report a per-amount intensive value of H, often normalized to a per-mole basis. mole of N2 and 1 mole of O2 is correct in this case because the standard enthalpy of formation always refers to 1 mole of product, NO2(g). And that means the combustion of ethanol is an exothermic reaction. the the bond enthalpies of the bonds broken. Some reactions are difficult, if not impossible, to investigate and make accurate measurements for experimentally. Step 2: Write out what you want to solve (eq. The one is referring to breaking one mole of carbon-carbon single bonds. with 348 kilojoules per mole for our calculation. Open Stax (examples and exercises). So this was 348 kilojoules per one mole of carbon-carbon single bonds. On the other hand, the heat produced by a reaction measured in a bomb calorimeter (Figure 5.17) is not equal to H because the closed, constant-volume metal container prevents the pressure from remaining constant (it may increase or decrease if the reaction yields increased or decreased amounts of gaseous species). Thus, the symbol (H)(H) is used to indicate an enthalpy change for a process occurring under these conditions. Given: Enthalpies of formation: C 2 H 5 O H ( l ), 278 kJ/mol. By signing up you are agreeing to receive emails according to our privacy policy. Here I just divided the 1354 by 2 to obtain the number of the energy released when one mole is burned. for the formation of C2H2). This finding (overall H for the reaction = sum of H values for reaction steps in the overall reaction) is true in general for chemical and physical processes. 2 Measure 100ml of water into the tin can. This ratio, (286kJ2molO3),(286kJ2molO3), can be used as a conversion factor to find the heat produced when 1 mole of O3(g) is formed, which is the enthalpy of formation for O3(g): Therefore, Hf[ O3(g) ]=+143 kJ/mol.Hf[ O3(g) ]=+143 kJ/mol. Figure \(\PageIndex{2}\): The steps of example \(\PageIndex{1}\) expressed as an energy cycle. Balance each of the following equations by writing the correct coefficient on the line. (credit: modification of work by AlexEagle/Flickr), Emerging Algae-Based Energy Technologies (Biofuels), (a) Tiny algal organisms can be (b) grown in large quantities and eventually (c) turned into a useful fuel such as biodiesel. To create this article, volunteer authors worked to edit and improve it over time. So the bond enthalpy for our carbon-oxygen double The LibreTexts libraries arePowered by NICE CXone Expertand are supported by the Department of Education Open Textbook Pilot Project, the UC Davis Office of the Provost, the UC Davis Library, the California State University Affordable Learning Solutions Program, and Merlot. (b) The density of ethanol is 0.7893 g/mL. Standard enthalpy of combustion (HC)(HC) is the enthalpy change when 1 mole of a substance burns (combines vigorously with oxygen) under standard state conditions; it is sometimes called heat of combustion. For example, the enthalpy of combustion of ethanol, 1366.8 kJ/mol, is the amount of heat produced when one mole of ethanol undergoes complete combustion at 25 C and 1 atmosphere pressure, yielding products also at 25 C and 1 atm. Microwave radiation has a wavelength on the order of 1.0 cm. This is described by the following equation, where where mi and ni are the stoichiometric coefficients of the products and reactants respectively. However, if we look For example, the molar enthalpy of formation of water is: \[H_2(g)+1/2O_2(g) \rightarrow H_2O(l) \; \; \Delta H_f^o = -285.8 \; kJ/mol \\ H_2(g)+1/2O_2(g) \rightarrow H_2O(g) \; \; \Delta H_f^o = -241.8 \; kJ/mol \]. So to get kilojoules as your final answer, if we go back up to here, we wrote a one times 348. It is only a rough estimate. And we're also not gonna worry And then for this ethanol molecule, we also have an The standard enthalpy change of the overall reaction is therefore equal to: (ii) the sum of the standard enthalpies of formation of all the products plus (i) the sum of the negatives of the standard enthalpies of formation of the reactants. This page titled 17.14: Heat of Combustion is shared under a CK-12 license and was authored, remixed, and/or curated by CK-12 Foundation via source content that was edited to the style and standards of the LibreTexts platform; a detailed edit history is available upon request. The next step is to look work is done on the system by the surroundings 10. And since it takes energy to break bonds, energy is given off when bonds form. This material has bothoriginal contributions, and contentbuilt upon prior contributions of the LibreTexts Community and other resources,including but not limited to: This page titled 5.7: Enthalpy Calculations is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by Robert Belford. The burning of ethanol produces a significant amount of heat. https://chem.libretexts.org/Bookshelves/Introductory_Chemistry/Book%3A_Introductory_Chemistry_(CK-12)/17%3A_Thermochemistry/17.14%3A_Heat_of_Combustion, https://courses.lumenlearning.com/boundless-chemistry/chapter/calorimetry/, https://sciencing.com/calculate-heat-absorption-6641786.html, https://chem.libretexts.org/Bookshelves/General_Chemistry/Book%3A_General_Chemistry_Supplement_(Eames)/Thermochemistry/Hess'_Law_and_Enthalpy_of_Formation, https://ch301.cm.utexas.edu/section2.php?target=thermo/thermochemistry/hess-law.html. This problem is solved in video \(\PageIndex{1}\) above. So we would need to break three The standard enthalpy of combustion is #H_"c"^#. When we add these together, we get 5,974. We can calculate the heating value using a steady-state energy balance on the stoichiometric reaction per 1 kmole of fuel, at constant temperature, and assuming complete combustion. For more tips, including how to calculate the heat of combustion with an experiment, read on. What are the units used for the ideal gas law? We also acknowledge previous National Science Foundation support under grant numbers 1246120, 1525057, and 1413739. oxygen-hydrogen single bonds. So we have one carbon-carbon bond. (credit a: modification of work by Micah Sittig; credit b: modification of work by Robert Kerton; credit c: modification of work by John F. Williams). And so, that's how to end up with kilojoules as your final answer. It produces somewhat lower carbon monoxide and carbon dioxide emissions, but does increase air pollution from other materials. Under the conditions of the reaction, methanol forms as a gas. Which of the following is an endothermic process? Calculate the enthalpy of combustion of exactly 1 L of ethanol. about units until the end, just to save some space on the screen. For example, energy is transferred into room-temperature metal wire if it is immersed in hot water (the wire absorbs heat from the water), or if you rapidly bend the wire back and forth (the wire becomes warmer because of the work done on it). citation tool such as, Authors: Paul Flowers, Klaus Theopold, Richard Langley, William R. Robinson, PhD. Bond enthalpies can be used to estimate the change in enthalpy for a chemical reaction. 7.!!4!g!of!acetylene!was!combusted!in!a!bomb!calorimeter!that!had!a!heat!capacity!of! In fact, it is not even a combustion reaction. An exothermic reaction is a reaction is which energy is given off to the surroundings, and enthalpy of reaction is the change in energy the atoms and molecules taking part in the reaction undergo. Science Chemistry Chemistry questions and answers Calculate the heat of combustion for one mole of acetylene (C2H2) using the following information. times the bond enthalpy of an oxygen-hydrogen single bond. To figure out which bonds are broken and which bonds are formed, it's helpful to look at the dot structures for our molecules. This calculator provides a way to compare the cost for various fuels types. And this now gives us the Enthalpy values for specific substances cannot be measured directly; only enthalpy changes for chemical or physical processes can be determined. \[\ce{N2}(g)+\ce{2O2}(g)\ce{2NO2}(g) \nonumber\], \[\ce{N2}(g)+\ce{O2}(g)\ce{2NO}(g)\hspace{20px}H=\mathrm{180.5\:kJ} \nonumber\], \[\ce{NO}(g)+\frac{1}{2}\ce{O2}(g)\ce{NO2}(g)\hspace{20px}H=\mathrm{57.06\:kJ} \nonumber\]. Since summing these three modified reactions yields the reaction of interest, summing the three modified H values will give the desired H: Aluminum chloride can be formed from its elements: (i) \(\ce{2Al}(s)+\ce{3Cl2}(g)\ce{2AlCl3}(s)\hspace{20px}H=\:?\), (ii) \(\ce{HCl}(g)\ce{HCl}(aq)\hspace{20px}H^\circ_{(ii)}=\mathrm{74.8\:kJ}\), (iii) \(\ce{H2}(g)+\ce{Cl2}(g)\ce{2HCl}(g)\hspace{20px}H^\circ_{(iii)}=\mathrm{185\:kJ}\), (iv) \(\ce{AlCl3}(aq)\ce{AlCl3}(s)\hspace{20px}H^\circ_{(iv)}=\mathrm{+323\:kJ/mol}\), (v) \(\ce{2Al}(s)+\ce{6HCl}(aq)\ce{2AlCl3}(aq)+\ce{3H2}(g)\hspace{20px}H^\circ_{(v)}=\mathrm{1049\:kJ}\). We did this problem, assuming that all of the bonds that we drew in our dots As we concentrate on thermochemistry in this chapter, we need to consider some widely used concepts of thermodynamics. Using Hesss Law Chlorine monofluoride can react with fluorine to form chlorine trifluoride: (i) \(\ce{ClF}(g)+\ce{F2}(g)\ce{ClF3}(g)\hspace{20px}H=\:?\). So next, we're gonna The heat of combustion of acetylene is -1309.5 kJ/mol. Research source. Explain how you can confidently determine the identity of the metal). Amount of ethanol used: \[\frac{1.55 \: \text{g}}{46.1 \: \text{g/mol}} = 0.0336 \: \text{mol}\nonumber \], Energy generated: \[4.184 \: \text{J/g}^\text{o} \text{C} \times 200 \: \text{g} \times 55^\text{o} \text{C} = 46024 \: \text{J} = 46.024 \: \text{kJ}\nonumber \], Molar heat of combustion: \[\frac{46.024 \: \text{kJ}}{0.0336 \: \text{mol}} = 1370 \: \text{kJ/mol}\nonumber \]. You can find these in a table from the CRC Handbook of Chemistry and Physics. Hesss law is valid because enthalpy is a state function: Enthalpy changes depend only on where a chemical process starts and ends, but not on the path it takes from start to finish. Both have the same change in elevation (altitude or elevation on a mountain is a state function; it does not depend on path), but they have very different distances traveled (distance walked is not a state function; it depends on the path). - [Educator] Bond enthalpies can be used to estimate the standard You will need to understand why it works..Hess Law states that the enthalpies of the products and the reactants are the same, All tip submissions are carefully reviewed before being published. \[\Delta H_{reaction}=\sum m_i \Delta H_{f}^{o}(products) - \sum n_i \Delta H_{f}^{o}(reactants) \nonumber \]. To get this, reverse and halve reaction (ii), which means that the H changes sign and is halved: To get ClF3 as a product, reverse (iv), changing the sign of H: Now check to make sure that these reactions add up to the reaction we want: Reactants 12O212O2 That is, the energy lost in the exothermic steps of the cycle must be regained in the endothermic steps, no matter what those steps are. This is usually rearranged slightly to be written as follows, with representing the sum of and n standing for the stoichiometric coefficients: The following example shows in detail why this equation is valid, and how to use it to calculate the enthalpy change for a reaction of interest. How do you calculate the ideal gas law constant? times the bond enthalpy of an oxygen-oxygen double bond. The heating value is then. Example \(\PageIndex{4}\): Writing Reaction Equations for \(H^\circ_\ce{f}\). Our goal is to manipulate and combine reactions (ii), (iii), and (iv) such that they add up to reaction (i). Table \(\PageIndex{1}\) Heats of combustion for some common substances. We're gonna approach this problem first like we're breaking all of Here, in the above reaction, one mole of acetylene produces -1301.1 kJ heat. You also might see kilojoules So looking at the ethanol molecule, we would need to break Molar enthalpies of formation are intensive properties and are the enthalpy per mole, that is the enthalpy change associated with the formation of one mole of a substance from its elements in their standard states. Do not include units in you answer C2H2 (g) + O2 (g) - 2C02 (g) + H20 (9) Bond C-C CEC Bond Energy (kJ/mol) 347 614 839 C-H C=0 O-H This problem has been solved! Looking at our balanced equation, we have one mole of ethanol reacting with three moles of oxygen gas to produce two moles of carbon dioxide and three moles of water to what we wrote here, we show breaking one oxygen-hydrogen A more comprehensive table can be found at the table of standard enthalpies of formation , which will open in a new window, and was taken from the CRC Handbook of Chemistry and Physics, 84 Edition (2004). carbon-oxygen single bond. For example, consider the following reaction phosphorous reacts with oxygen to from diphosphorous pentoxide (2P2O5), \[P_4+5O_2 \rightarrow 2P_2O_5\] This calculator provides a quick way to compare the cost and CO2 emissions for various fuels. Before we further practice using Hesss law, let us recall two important features of H. negative sign in here because this energy is given off. carbon-oxygen double bonds. times the bond enthalpy of a carbon-oxygen double bond. Paul Flowers, Klaus Theopold, Richard Langley, (c) Calculate the heat of combustion of 1 mole of liquid methanol to H. Chemists usually perform experiments under normal atmospheric conditions, at constant external pressure with q = H, which makes enthalpy the most convenient choice for determining heat changes for chemical reactions. Robert E. Belford (University of Arkansas Little Rock; Department of Chemistry). Specific heat capacity is the quantity of heat needed to change the temperature of 1.00 g of a substance by 1 K. 11. 1molrxn 1molC 2 H 2)(1molC 2 H 26gC 2 H 2)(4gC 2 H 2) H 4g =200kJ U=q+w U 4g =200,000J+571.7J=199.4kJ!!! \nonumber\]. How much heat is produced by the combustion of 125 g of acetylene? This leaves only reactants ClF(g) and F2(g) and product ClF3(g), which are what we want. How much heat is produced by the combustion of 125 g of acetylene? Finally, let's show how we get our units. !What!is!the!expected!temperature!change!in!such!a . Calculate the heat evolved/absorbed given the masses (or volumes) of reactants. of reaction as our units, the balanced equation had Start by writing the balanced equation of combustion of the substance. The trick is to add the above equations to produce the equation you want. \[\begin{align} \cancel{\color{red}{2CO_2(g)}} + \cancel{\color{green}{H_2O(l)}} \rightarrow C_2H_2(g) +\cancel{\color{blue} {5/2O_2(g)}} \; \; \; \; \; \; & \Delta H_{comb} = -(-\frac{-2600kJ}{2} ) \nonumber \\ \nonumber \\ 2C(s) + \cancel{\color{blue} {2O_2(g)}} \rightarrow \cancel{\color{red}{2CO_2(g)}} \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; & \Delta H_{comb}= 2(-393 kJ) \nonumber \\ \nonumber \\ H_2(g) +\cancel{\color{blue} {1/2O_2(g)}} \rightarrow \cancel{\color{green}{H_2O(l)}} \; \; \; \; \; \; \; \; \; \; \; & \Delta H_{comb} = \frac{-572kJ}{2} \end{align}\], Step 4: Sum the Enthalpies: 226kJ (the value in the standard thermodynamic tables is 227kJ, which is the uncertain digit of this number). We can look at this as a two step process. to sum the bond enthalpies of the bonds that are formed. See Answer 447 kJ B. The stepwise reactions we consider are: (i) decompositions of the reactants into their component elements (for which the enthalpy changes are proportional to the negative of the enthalpies of formation of the reactants), followed by (ii) re-combinations of the elements to give the products (with the enthalpy changes proportional to the enthalpies of formation of the products). You can specify conditions of storing and accessing cookies in your browser. Calculate Hfor acetylene. Watch the video below to get the tips on how to approach this problem. Heats of combustion are usually determined by burning a known amount of the material in a bomb calorimeter with an excess of oxygen. The LibreTexts libraries arePowered by NICE CXone Expertand are supported by the Department of Education Open Textbook Pilot Project, the UC Davis Office of the Provost, the UC Davis Library, the California State University Affordable Learning Solutions Program, and Merlot. Click here to learn more about the process of creating algae biofuel. up the bond enthalpies of all of these different bonds. Subtract the initial temperature of the water from 40 C. Substitute it into the formula and you will get the answer q in J. The heat of combustion of. If the sum of the bond enthalpies of the bonds that are broken, if this number is larger than the sum of the bond enthalpies of the bonds that have formed, we would've gotten a positive value for the change in enthalpy. As such, enthalpy has the units of energy (typically J or cal). From data tables find equations that have all the reactants and products in them for which you have enthalpies. The enthalpy of formation, \(H^\circ_\ce{f}\), of FeCl3(s) is 399.5 kJ/mol. So the summation of the bond enthalpies of the bonds that are broken is going to be a positive value. Our mission is to improve educational access and learning for everyone. The breadth, depth and veracity of this work is the responsibility of Robert E. Belford, rebelford@ualr.edu. If so how is a negative enthalpy indicate an exothermic reaction? If we scrutinise this statement: "the total energies of the products being less than the reactants", then a negative enthalpy cannot be an exothermic. Direct link to Morteza Aslami's post what do we mean by bond e, Posted a month ago. an endothermic reaction. using the above equation, we get, Example \(\PageIndex{3}\) Calculating enthalpy of reaction with hess's law and combustion table, Using table \(\PageIndex{1}\) Calculate the enthalpy of reaction for the hydrogenation of ethene into ethane, \[C_2H_4 + H_2 \rightarrow C_2H_6 \nonumber \]. Also notice that the sum 265897 views And since we're For the reaction H2(g)+Cl2(g)2HCl(g)H=184.6kJH2(g)+Cl2(g)2HCl(g)H=184.6kJ, (a) 2C(s,graphite)+3H2(g)+12O2(g)C2H5OH(l)2C(s,graphite)+3H2(g)+12O2(g)C2H5OH(l), (b) 3Ca(s)+12P4(s)+4O2(g)Ca3(PO4)2(s)3Ca(s)+12P4(s)+4O2(g)Ca3(PO4)2(s). bond is 799 kilojoules per mole, and we multiply that by four. That is, you can have half a mole (but you can not have half a molecule. OpenStax is part of Rice University, which is a 501(c)(3) nonprofit. Next, we do the same thing for the bond enthalpies of the bonds that are formed. How does Charle's law relate to breathing? Base heat released on complete consumption of limiting reagent. To create this article, volunteer authors worked to edit and improve it over time. H 2 O ( l ), 286 kJ/mol. We will consider how to determine the amount of work involved in a chemical or physical change in the chapter on thermodynamics. So we write a one, and then the bond enthalpy for a carbon-oxygen single bond. (The symbol H is used to indicate an enthalpy change for a reaction occurring under nonstandard conditions. The calculator takes into account the cost of the fuel, energy content of the fuel, and the efficiency of your furnace. The answer is the experimental heat of combustion in kJ/g. The heat of combustion refers to the energy that is released as heat when a compound undergoes complete combustion with oxygen under standard conditions. Except where otherwise noted, textbooks on this site If 1 mol of acetylene produces -1301.1 kJ, then 4.8 mol of acetylene produces: \(\begin{array}{l}{\rm{ = 1301}}{\rm{.1 \times 4}}{\rm{.8 }}\\{\rm{ = 6245}}{\rm{.28 kJ }}\\{\rm{ = 6}}{\rm{.25 kJ}}\end{array}\). Textbook content produced by OpenStax is licensed under a Creative Commons Attribution License . Going from left to right in (i), we first see that \(\ce{ClF}_{(g)}\) is needed as a reactant. 94% of StudySmarter users get better grades. We recommend using a However, we're gonna go Determine the heat released or absorbed when 15.0g Al react with 30.0g Fe3O4(s). So, identify species that only exist in one of the given equations and put them on the desired side of the equation you want to produce, following the Tips above. The work, w, is positive if it is done on the system and negative if it is done by the system. Summing these reaction equations gives the reaction we are interested in: Summing their enthalpy changes gives the value we want to determine: So the standard enthalpy change for this reaction is H = 138.4 kJ. five times the bond enthalpy of an oxygen-hydrogen single bond. Calculate the molar heat of combustion. By their definitions, the arithmetic signs of V and w will always be opposite: Substituting this equation and the definition of internal energy into the enthalpy-change equation yields: where qp is the heat of reaction under conditions of constant pressure. Accessibility StatementFor more information contact us atinfo@libretexts.orgor check out our status page at https://status.libretexts.org. In this class, the standard state is 1 bar and 25C. Its energy contentis H o combustion = -1212.8kcal/mole. Do the same for the reactants. Posted 2 years ago. Note: If you do this calculation one step at a time, you would find: As reserves of fossil fuels diminish and become more costly to extract, the search is ongoing for replacement fuel sources for the future. See video \(\PageIndex{2}\) for tips and assistance in solving this.