estimate the heat of combustion for one mole of acetylene

For the formation of 2 mol of O3(g), H=+286 kJ.H=+286 kJ. 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. !What!is!the!expected!temperature!change!in!such!a . When you multiply these two together, the moles of carbon-carbon Specific heat capacity is the quantity of heat needed to change the temperature of 1.00 g of a substance by 1 K. 11. Find the amount of substance burned by subtracting the final mass from the initial mass of the substance in g. Divide q in kJ by the mass of the substance burned. then you must include on every physical page the following attribution: If you are redistributing all or part of this book in a digital format, And we continue with everything else for the summation of From data tables find equations that have all the reactants and products in them for which you have enthalpies. Enthalpies of formation are usually found in a table from CRC Handbook of Chemistry and Physics. By using the following special form of the Hess' law, we can calculate the heat of combustion of 1 mole of ethanol. (The symbol H is used to indicate an enthalpy change for a reaction occurring under nonstandard conditions. This is also the procedure in using the general equation, as shown. Calculate the heat of combustion . The chemical reaction is given in the equation; The bond energy of the reactant is: Following the bond energies given in the question, we have: = ( 1 839) + (5/2 495) + (2 413) The species of algae used are nontoxic, biodegradable, and among the worlds fastest growing organisms. 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. wikiHow is a wiki, similar to Wikipedia, which means that many of our articles are co-written by multiple authors. Example \(\PageIndex{4}\): Writing Reaction Equations for \(H^\circ_\ce{f}\). Substances act as reservoirs of energy, meaning that energy can be added to them or removed from them. You should contact him if you have any concerns. So let's write in here, the bond enthalpy for An example of a state function is altitude or elevation. Determine the heat released or absorbed when 15.0g Al react with 30.0g Fe3O4(s). Sign up for free to discover our expert answers. Before we further practice using Hesss law, let us recall two important features of H. The standard enthalpy of formation of CO2(g) is 393.5 kJ/mol. 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. Convert into kJ by dividing q by 1000. Next, we have to break a structures were formed. Determine the specific heat and the identity of the metal. \[30.0gFe_{3}O_{4}\left(\frac{1molFe_{3}O_{4}}{231.54g}\right) \left(\frac{-3363kJ}{3molFe_{3}O_{4}}\right) = -145kJ\], Note, you could have used the 0.043 from step 2, Thanks to all authors for creating a page that has been read 135,840 times. This book uses the Direct link to JPOgle 's post An exothermic reaction is. Watch Video \(\PageIndex{1}\) to see these steps put into action while solving example \(\PageIndex{1}\). So to this, we're going to write in here, a five, and then the bond enthalpy of a carbon-hydrogen bond. And the 348, of course, is the bond enthalpy for a carbon-carbon single bond. Let's use bond enthalpies to estimate the enthalpy of combustion of ethanol. However, if we look 265897 views Research source. The calculator estimates the cost for each fuel type to deliver 100,000 BTU's of heat to your house. Solved Estimate the heat of combustion for one mole of - Chegg This can be obtained by multiplying reaction (iii) by \(\frac{1}{2}\), which means that the H change is also multiplied by \(\frac{1}{2}\): \[\ce{ClF}(g)+\frac{1}{2}\ce{O2}(g)\frac{1}{2}\ce{Cl2O}(g)+\frac{1}{2}\ce{OF2}(g)\hspace{20px} H=\frac{1}{2}(205.6)=+102.8\: \ce{kJ} \nonumber\]. When we add these together, we get 5,974. 6.7: Tabulated Enthalpy Values - Chemistry LibreTexts 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. (This amount of energy is enough to melt 99.2 kg, or about 218 lbs, of ice.) If you're behind a web filter, please make sure that the domains *.kastatic.org and *.kasandbox.org are unblocked. 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: \[\begin {align*} 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). (Figure 6 in Chapter 5.1 Energy Basics) is essentially pure acetylene, the heat produced by combustion of one mole of acetylene in such a torch is likely not equal to the enthalpy of combustion of acetylene listed in Table 2. Free and expert-verified textbook solutions. Since summing these three modified reactions yields the reaction of interest, summing the three modified H values will give the desired H: (i) 2Al(s)+3Cl2(g)2AlCl3(s)H=?2Al(s)+3Cl2(g)2AlCl3(s)H=? Ethanol, C 2 H 5 OH, is used as a fuel for motor vehicles, particularly in Brazil. of energy are given off for the combustion of one mole of ethanol. We can look at this in an Energy Cycle Diagram (Figure \(\PageIndex{2}\)). This "gasohol" is widely used in many countries. The calculator estimates the cost and CO2 emissions for each fuel to deliver 100,000 BTU's of heat to your house. The heat given off when you operate a Bunsen burner is equal to the enthalpy change of the methane combustion reaction that takes place, since it occurs at the essentially constant pressure of the atmosphere. The bonds enthalpy for an oxygen hydrogen single bond is 463 kilojoules per mole, and we multiply that by six. Figure \(\PageIndex{2}\): The steps of example \(\PageIndex{1}\) expressed as an energy cycle. ), The enthalpy changes for many types of chemical and physical processes are available in the reference literature, including those for combustion reactions, phase transitions, and formation reactions. work is done on the system by the surroundings 10. \[30.0gFe_{3}O_{4}\left(\frac{1molFe_{3}O_{4}}{231.54g}\right) \left(\frac{1}{3molFe_{3}O_{4}}\right) = 0.043\], From T1: Standard Thermodynamic Quantities we obtain the enthalpies of formation, Hreaction = mi Hfo (products) ni Hfo (reactants), Hreaction = 4(-1675.7) + 9(0) -8(0) -3(-1118.4)= -3363.6kJ. This view of an internal combustion engine illustrates the conversion of energy produced by the exothermic combustion reaction of a fuel such as gasoline into energy of motion. Using Hesss Law Determine the enthalpy of formation, \(H^\circ_\ce{f}\), of FeCl3(s) from the enthalpy changes of the following two-step process that occurs under standard state conditions: \[\ce{Fe}(s)+\ce{Cl2}(g)\ce{FeCl2}(s)\hspace{20px}H=\mathrm{341.8\:kJ} \nonumber\], \[\ce{FeCl2}(s)+\frac{1}{2}\ce{Cl2}(g)\ce{FeCl3}(s)\hspace{20px}H=\mathrm \nonumber{57.7\:kJ} \]. Here is a video that discusses how to calculate the enthalpy change when 0.13 g of butane is burned. (b) The first time a student solved this problem she got an answer of 88 C. \end {align*}\]. a carbon-carbon bond. Solved Calculate the heat of combustion for one mole of | Chegg.com And, kilojoules per mole reaction means how the reaction is written. 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. We use cookies to make wikiHow great. So let's start with the ethanol molecule. \[\begin{align} 2C_2H_2(g) + 5O_2(g) \rightarrow 4CO_2(g) + 2H_2O(l) \; \; \; \; \; \; & \Delta H_{comb} =-2600kJ \nonumber \\ C(s) + O_2(g) \rightarrow CO_2(g) \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; & \Delta H_{comb}= -393kJ \nonumber \\ 2H_2(g) + O_2 \rightarrow 2H_2O(l) \; \; \; \; \; \; \; \; \; \; \; \;\; \; \; \; \; \; & \Delta H_{comb} = -572kJ \end{align}\]. If so how is a negative enthalpy indicate an exothermic reaction? 3.51kJ/Cforthedevice andcontained2000gofwater(C=4.184J/ g!C)toabsorb! Also, these are not reaction enthalpies in the context of a chemical equation (section 5.5.2), but the energy per mol of substance combusted. It takes energy to break a bond. To calculate the heat of combustion, use Hesss law, which states that the enthalpies of the products and the reactants are the same. sum of the bond enthalpies for all the bonds that need to be broken. subtracting a larger number from a smaller number, we get that negative sign for the change in enthalpy. 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. Use the formula q = Cp * m * (delta) t to calculate the heat liberated which heats the water. Our goal is to manipulate and combine reactions (ii), (iii), and (iv) such that they add up to reaction (i). Write the equation you want on the top of your paper, and draw a line under it. The heat of combustion refers to the energy that is released as heat when a compound undergoes complete combustion with oxygen under standard conditions. The distances traveled would differ (distance is not a state function) but the elevation reached would be the same (altitude is a state function). 27 febrero, 2023 . For each product, you multiply its #H_"f"^# by its coefficient in the balanced equation and add them together. That is, the equation in the video and the one above have the exact same value, just one is per mole, the other is per 2 mols of acetylene. calculate the number of N, C, O, and H atoms in 1.78*10^4g of urea. This way it is easier to do dimensional analysis. The next step is to look \nonumber\]. times the bond enthalpy of an oxygen-oxygen double bond. Now, when we multiply through the moles of carbon-carbon single bonds, cancel and this gives us Fuel Comparison Calculator - Build-It-Solar (The engine is able to keep the car moving because this process is repeated many times per second while the engine is running.) Dec 15, 2022 OpenStax. How to calculate the heat released by the combustion of ethanol in (ii) HCl(g)HCl(aq)H(ii)=74.8kJHCl(g)HCl(aq)H(ii)=74.8kJ, (iii) H2(g)+Cl2(g)2HCl(g)H(iii)=185kJH2(g)+Cl2(g)2HCl(g)H(iii)=185kJ, (iv) AlCl3(aq)AlCl3(s)H(iv)=+323kJ/molAlCl3(aq)AlCl3(s)H(iv)=+323kJ/mol, (v) 2Al(s)+6HCl(aq)2AlCl3(aq)+3H2(g)H(v)=1049kJ2Al(s)+6HCl(aq)2AlCl3(aq)+3H2(g)H(v)=1049kJ. If the equation has a different stoichiometric coefficient than the one you want, multiply everything by the number to make it what you want, including the reaction enthalpy, \(\Delta H_2\) = -1411kJ/mol Total Exothermic = -1697 kJ/mol, \(\Delta H_4\) = - \(\Delta H^*_{rxn}\) = ? You'll get a detailed solution from a subject matter expert that helps you learn core concepts. In our balanced equation, we formed two moles of carbon dioxide. Note, these are negative because combustion is an exothermic reaction. Explain how you can confidently determine the identity of the metal). In this case, one mole of oxygen reacts with one mole of methanol to form one mole of carbon dioxide and two moles of water. \[\Delta H_1 +\Delta H_2 + \Delta H_3 + \Delta H_4 = 0\]. The heat(enthalpy) of combustion of acetylene = -1228 kJ. Calculate the sodium ion concentration when 70.0 mL of 3.0 M sodium carbonate is added to 30.0 mL of 1.0 M sodium bicarbonate. Subtract the reactant sum from the product sum. \[\Delta H_{reaction}=\sum m_i \Delta H_{f}^{o}(products) - \sum n_i \Delta H_{f}^{o}(reactants) \\ where \; m_i \; and \; n_i \; \text{are the stoichiometric coefficients of the products and reactants respectively} \]. 94% of StudySmarter users get better grades. Accessibility StatementFor more information contact us atinfo@libretexts.orgor check out our status page at https://status.libretexts.org. For example, given that: Then, for the reverse reaction, the enthalpy change is also reversed: Looking at the reactions, we see that the reaction for which we want to find H is the sum of the two reactions with known H values, so we must sum their Hs: The enthalpy of formation, Hf,Hf, of FeCl3(s) is 399.5 kJ/mol. The number of moles of acetylene is calculated as: Learn more about heat of combustion here: This site is using cookies under cookie policy . This article has been viewed 135,840 times. We also acknowledge previous National Science Foundation support under grant numbers 1246120, 1525057, and 1413739. #DeltaH_("C"_2"H"_2"(g)")^o = "226.73 kJ/mol"#; #DeltaH_("CO"_2"(g)")^o = "-393.5 kJ/mol"#; #DeltaH_("H"_2"O(l)")^o = "-285.8 kJ/mol"#, #"[2 (-393.5) + (-295.8)] [226.7 + 0] kJ" = "-1082.8 - 226.7" =#. a carbon-carbon bond. If we look at the process diagram in Figure \(\PageIndex{3}\) and correlate it to the above equation we see two things. So next, we're gonna 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. around the world. single bonds over here, and we show the formation of six oxygen-hydrogen This problem is solved in video \(\PageIndex{1}\) above. And we can see that in To log in and use all the features of Khan Academy, please enable JavaScript in your browser. Calculate the heat of combustion for one mole of acetylene. - OneClass single bonds over here. write this down here. If gaseous water forms, only 242 kJ of heat are released. 125 g of acetylene produces 6.25 kJ of heat. How much heat is produced by the combustion of 125 g of acetylene? The provided amounts of the two reactants are, The provided molar ratio of perchlorate-to-sucrose is then. Ethanol (CH 3 CH 2 OH) has H o combustion = -326.7 kcal/mole. consent of Rice University. They are often tabulated as positive, and it is assumed you know they are exothermic. for the formation of C2H2). 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. A standard enthalpy of formation HfHf is an enthalpy change for a reaction in which exactly 1 mole of a pure substance is formed from free elements in their most stable states under standard state conditions. of the area used to grow corn) can produce enough algal fuel to replace all the petroleum-based fuel used in the US. So if you look at your dot structures, if you see a bond that's the Using the following bond energies: Bond Bond Energy (kJ/mol) - BRAINLY bond is about 348 kilojoules per mole. Notice that we got a negative value for the change in enthalpy. And since we have three moles, we have a total of six water that's drawn here, we form two oxygen-hydrogen single bonds. A 1.55 gram sample of ethanol is burned and produced a temperature increase of \(55^\text{o} \text{C}\) in 200 grams of water. https://openstax.org/books/chemistry-2e/pages/1-introduction, https://openstax.org/books/chemistry-2e/pages/5-3-enthalpy, Creative Commons Attribution 4.0 International License, Define enthalpy and explain its classification as a state function, Write and balance thermochemical equations, Calculate enthalpy changes for various chemical reactions, Explain Hesss law and use it to compute reaction enthalpies. Kilimanjaro. According to the US Department of Energy, only 39,000 square kilometers (about 0.4% of the land mass of the US or less than 1717 single bonds cancels and this gives you 348 kilojoules. For nitrogen dioxide, NO2(g), HfHf is 33.2 kJ/mol. Calculate the heat evolved/absorbed given the masses (or volumes) of reactants. In this case, there is no water and no carbon dioxide formed. We will include a superscripted o in the enthalpy change symbol to designate standard state. Posted 2 years ago. bond is 799 kilojoules per mole, and we multiply that by four. Amount of ethanol used: 1.55 g 46.1 g/mol = 0.0336 mol Energy generated: Does it mean the amount of energies required to break or form bonds? This type of calculation usually involves the use of Hesss law, which states: If a process can be written as the sum of several stepwise processes, the enthalpy change of the total process equals the sum of the enthalpy changes of the various steps. The heating value is then. You might see a different value, if you look in a different textbook. \[\begin{align} \cancel{\color{red}{2CO_2(g)}} + \cancel{\color{green}{H_2O(l)}} \rightarrow C_2H_2(g) +\cancel{\color{blue} {5/2O_2(g)}} \; \; \; \; \; \; & \Delta H_{comb} = -(-\frac{-2600kJ}{2} ) \nonumber \\ \nonumber \\ 2C(s) + \cancel{\color{blue} {2O_2(g)}} \rightarrow \cancel{\color{red}{2CO_2(g)}} \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; \; & \Delta H_{comb}= 2(-393 kJ) \nonumber \\ \nonumber \\ H_2(g) +\cancel{\color{blue} {1/2O_2(g)}} \rightarrow \cancel{\color{green}{H_2O(l)}} \; \; \; \; \; \; \; \; \; \; \; & \Delta H_{comb} = \frac{-572kJ}{2} \end{align}\], Step 4: Sum the Enthalpies: 226kJ (the value in the standard thermodynamic tables is 227kJ, which is the uncertain digit of this number). If you're seeing this message, it means we're having trouble loading external resources on our website. times the bond enthalpy of an oxygen-hydrogen single bond. References. If the coefficients of the chemical equation are multiplied by some factor, the enthalpy change must be multiplied by that same factor (H is an extensive property): The enthalpy change of a reaction depends on the physical states of the reactants and products, so these must be shown. Many thermochemical tables list values with a standard state of 1 atm. This calculator provides a way to compare the cost for various fuels types. Calculate the molar heat of combustion. Click here to learn more about the process of creating algae biofuel. Enthalpies of combustion for many substances have been measured; a few of these are listed in Table 5.2. How do you calculate the ideal gas law constant? Under the conditions of the reaction, methanol forms as a gas. It is often important to know the energy produced in such a reaction so that we can determine which fuel might be the most efficient for a given purpose. Algae can yield 26,000 gallons of biofuel per hectaremuch more energy per acre than other crops. Question. 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. 7.!!4!g!of!acetylene!was!combusted!in!a!bomb!calorimeter!that!had!a!heat!capacity!of! Many chemical reactions are combustion reactions. Table \(\PageIndex{1}\) Heats of combustion for some common substances. For example, C2H2(g) + 5 2O2(g) 2CO2(g) +H2O (l) You calculate H c from standard enthalpies of formation: H o c = H f (p) H f (r) So to represent the three And 1,255 kilojoules The Heat of Combustion of a substance is defined as the amount of energy in the form of heat is liberated when an amount of the substance undergoes combustion. 0.043(-3363kJ)=-145kJ. wikiHow is a wiki, similar to Wikipedia, which means that many of our articles are co-written by multiple authors. 5.3 Enthalpy - Chemistry For example, we can think of the reaction of carbon with oxygen to form carbon dioxide as occurring either directly or by a two-step process. So looking at the ethanol molecule, we would need to break look at 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) \(\ce{2Al}(s)+\ce{6HCl}(aq)\ce{2AlCl3}(aq)+\ce{3H2}(g)\hspace{20px}H^\circ_{(v)}=\mathrm{1049\:kJ}\). Acetylene torches utilize the following reaction: 2 C2H2 (g 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. If an equation has a chemical on the opposite side, write it backwards and change the sign of the reaction enthalpy. Calculate the enthalpy of formation for acetylene, C2H2(g) from the combustion data (table \(\PageIndex{1}\), note acetylene is not on the table) and then compare your answer to the value in table \(\PageIndex{2}\), Hcomb (C2H2(g)) = -1300kJ/mol PDF Thermodynamics.Unit.1.RAQ. - University of Texas at Austin 2 See answers Advertisement Advertisement . Calculations using the molar heat of combustion are described. Legal. So to this, we're going to add a three Next, subtract the enthalpies of the reactants from the product. times the bond enthalpy of a carbon-oxygen double bond. - [Educator] Bond enthalpies can be used to estimate the standard (b) Methanol, a liquid fuel that could possibly replace gasoline, can be prepared from water gas and additional hydrogen at high temperature and pressure in the presence of a suitable catalyst:\({\bf{2}}{{\bf{H}}_{\bf{2}}}\left( {\bf{g}} \right){\bf{ + CO}}\left( {\bf{g}} \right) \to {\bf{C}}{{\bf{H}}_{\bf{3}}}{\bf{OH}}\left( {\bf{g}} \right)\). Base heat released on complete consumption of limiting reagent. what do we mean by bond enthalpies of bonds formed or broken? how much heat is produced by the combustion of 125 g of acetylene c2h2 Among the most promising biofuels are those derived from algae (Figure 5.22). &\ce{ClF}(g)+\frac{1}{2}\ce{O2}(g)\frac{1}{2}\ce{Cl2O}(g)+\frac{1}{2}\ce{OF2}(g)&&H=\mathrm{+102.8\: kJ}\\ The enthalpy of formation, \(H^\circ_\ce{f}\), of FeCl3(s) is 399.5 kJ/mol. Determine the total energy change for the production of one mole of aqueous nitric acid by this process. (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. Enthalpy is a state function which means the energy change between two states is independent of the path. The OpenStax name, OpenStax logo, OpenStax book covers, OpenStax CNX name, and OpenStax CNX logo Chemists use a thermochemical equation to represent the changes in both matter and energy. oxygen hydrogen single bond is 463 kilojoules per mole, and we multiply that by six. Enthalpy values for specific substances cannot be measured directly; only enthalpy changes for chemical or physical processes can be determined. And that would be true for [1] and 12O212O2 A 92.9-g piece of a silver/gray metal is heated to 178.0 C, and then quickly transferred into 75.0 mL of water initially at 24.0 C. Heats of combustion are usually determined by burning a known amount of the material in a bomb calorimeter with an excess of oxygen. So for the combustion of one mole of ethanol, 1,255 kilojoules of energy are released. negative sign in here because this energy is given off. In this video, we'll use average bond enthalpies to calculate the enthalpy change for the gas-phase combustion of ethanol. The combustion of 1.00 L of isooctane produces 33,100 kJ of heat. From table \(\PageIndex{1}\) we obtain the following enthalpies of combustion, \[\begin{align} \text{eq. If 1 mol of acetylene produces -1301.1 kJ, then 4.8 mol of acetylene produces: \(\begin{array}{l}{\rm{ = 1301}}{\rm{.1 \times 4}}{\rm{.8 }}\\{\rm{ = 6245}}{\rm{.28 kJ }}\\{\rm{ = 6}}{\rm{.25 kJ}}\end{array}\). Note, step 4 shows C2H6 -- > C2H4 +H2 and in example \(\PageIndex{1}\) we are solving for C2H4 +H2 --> C2H6 which is the reaction of step 4 written backwards, so the answer to \(\PageIndex{1}\) is the negative of step 4. the the bond enthalpies of the bonds broken. If a quantity is not a state function, then its value does depend on how the state is reached. By applying Hess's Law, H = H 1 + H 2.
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