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Practice: Hydrogenation of Acetylene

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Wize High School Grade 12 Chemistry Textbook > Energy Changes

Ways to Calculate Enthalpy of Reactions

10 Activities

Using the table below, answer the following questions about the hydrogenation reactions of acetylene:

a) Calculate the enthalpy of formation of ethylene, C2H4, considering ΔHrxn = -174.4kJ/mol
C2H2+H2→C2H4C_2H_2+H_2\to C_2H_4
Report your answer in kJ/mol to one decimal place. Do not include units in the answer field

b) Calculate So of H2(g) if ΔSrxn = -103.23 J/mol K for the reaction below
C2H4+H2→C2H6C_2H_4+H_2\to C_2H_6
Report your answer in J/mol K to one decimal place. Do not include units in the answer field

c) At what temperature does the first reaction above become spontaneous?

Report your answer in K with no decimal places. Do not include units in the answer field
More Ways to Calculate Enthalpy of Reactions Questions:
Enthalpy of Formation
What is the enthalpy of the following reaction when 0.35 g of Mg2+ is produced? Use the following information:
ΔH°f[Mg(OH)2(s)] = −924.5 kJ
ΔH°f[NH4+(aq)] = −132.5 kJ
Gibbs Free Energy
The reaction of
H2(g)+CO2(g)↔H2O(g)+CO(g) at 2000K KP=4.40H_{2(g)}+CO_{2(g)}\leftrightarrow H_2O_{(g)}+CO_{(g)}\ at\ 2000K\ K_P=4.40
Calculate ΔG for this reaction
Thermochemistry: Enthalpy of reactions
The following reaction enthalpies are provided:
H2(g)+12O2(g)→H2O(l),ΔH1=−285.8kJ/mol(1)H_{2(g)}+\frac{1}{2}O_{2(g)}\to H_2O_{(l)},\Delta H_1=-285.8kJ/mol\hspace{70pt} (1)
C(s)+O2(g)→CO2(g),ΔH2=−293.5kJ/mol(2)C_{(s)}+O_{2(g)}\to CO_{2(g)},\Delta H_2=-293.5kJ/mol \hspace{80pt} (2)
Calorimetry: Temperature change
1. A 56.0 g block of iron heated to 90 °C is placed in a 4.8 kg water bath at 20 °C. What will be the final temperature of the iron? heat capacity of iron c = 0.450 J /gK, heat capacity of water is 4.185 J/gK
Enthalpy of Reactions
Consider the following reaction:
N2H4(I)+O2→N2(g)+2 H2O(I)ΔH=−622.2 kJN_2H_4(I)+O_2\rightarrow N_2(g)+2\ H_2O(I)\quad \Delta H=-622.2\ kJ
Given the following data, calculate the heat of reaction for the same reaction where water is a gaseous product instead of a liquid:
Calculate ΔH°\Delta H\degree for the reaction below, which requires catalytic acid to proceed
C6H12(l)+H2O(l)→C6H13OH(l)C_6H_{12}\left(l\right)+H_2O\left(l\right)\to C_6H_{13}OH\left(l\right)
Given:
The standard enthalpy of formation for NO2(g)NO_2\left(g\right) is 33.18 kJ/mol, and the standard enthalpy of partial combustion of nitrogen to NO(g) is 90.25 kJ/mol. Calculate the standard enthalpy of combustion of NO(g) to NO2(g).
Hess’s Law: Indirect Determination of ΔH
The standard enthalpy of formation of ethanol (C2H5OH){\small \left(\text{C}_2\text{H}_5\text{OH}\right)} is -277.7 kJ/mol. Write the chemical equation to which this value applies.
Solution:
Calculate the ΔH for the reaction below IF5 (g) → IF3 (g) + F2 (g)
IF (g) + F2 (g) → IF3 (g) ΔH = -390 kJ IF (g) + 2 F2 (g) → IF5 (g) ΔH = -745 kJ
Liquification of fuels is an important process that allows fuel to be transported more safely and more easily around the world. Oxidizing methane transforms it from a gas to a liquid, methanol, which is much easier to transport. Using the data provided calculate ΔH\Delta Hfor the oxidation of methane (CH4) shown below
CH4(g) + 1/2 O2(g)→CH3OH(l)ΔHfo(CH4(g))=−74.9kJ mol−1ΔHfo(CH3OH(l)=−238.6 kJ mol−1\begin{array}{c} CH_{4(g)}\ +\ 1/2\ O_{2(g)} \rightarrow CH_3OH_{(l)} \\ \\ \Delta H^o_f(CH_{4(g)})=-74.9 kJ\ mol^{-1} \\ \Delta H_f^o(CH_3OH_{(l)}=-238.6\ kJ\ mol^{-1} \end{array}
Practice: Total Bond Energy of IF$_7$
The total bond enthalpy for F2 is 156.9 kJ/mol and the total bond enthalpy of I2 is 152.5 kJ/mol. Using the thermochemical equation below, what is the total bond enthalpy of IF7?
I2(s)+7 F2(g)→2 IF7(s)  ΔH=2669.2 kJ.molI_{2(s)}+7\ F_{2(g)}\rightarrow 2\ IF_{7(s)}\ \ \Delta H=2669.2\ kJ.mol
Enthalpy of Reactions
Given the following at 25°C and 1.00 atm:
ReactionΔH°12N2(g)+O2(g)→NO2(g)33.2kJN2(g)+2O2(g)→N2O4(g)11.1kJ\def\arraystretch{1.5}\begin{array}{cc}\rm Reaction&\Delta H\degree\\\cfrac{1}{2} N_2 (g)+O_2 (g)→NO_2 (g)&33.2 kJ \\N_2 (g)+2O_2 (g)→N_2 O_4 (g)&11.1 kJ\end{array}
Enthalpy of Reactions
Estimate the heat of reaction at 298 K for the reaction shown, given the average bond energies below.
Br2(g)+3F2(g)→2BrF3(g)Br_2 (g)+3F_2 (g)→2BrF_3 (g)
BondBond EnergyBr−Br193kJ/molF−F155kJ/molBr−F249kJ/mol\def\arraystretch{1.5} \begin{array}{c} \rm Bond &&&& \rm Bond \ Energy \\ \hline Br-Br &&&& 193 kJ/mol\\ F-F &&&& 155 kJ/mol\\ Br-F &&&& 249 kJ/mol\\ \end{array}
Enthalpy: Bond Dissociation Energies
Consider the following balanced chemical reaction and the bond dissociation energies for the partial combustion of hydrogen sulfide:
What is the enthalpy change ΔrH for this reaction?
Thermochemistry: Calorimetry
Consider the following balanced chemical equation:
2CℓF3(g)+2O2(g)→Cℓ2O(g)+3F2O(g)∆rH°=???2 CℓF_3(g) + 2 O_2(g) → Cℓ_2O (g) + 3 F_2O (g) \qquad ∆_rH° = ???
Substance△fH° @ 25°C (kJ/mol)CℓF3(g)??Cℓ2O(g)80.3F2O(g)–21.7F2(g)0\def\arraystretch{1.5} \begin{array}{|c|c|} \hline \rm Substance & △_fH° \ @ \ 25°C \ (kJ/mol) \\ \hline CℓF_3 (g) & ?? \\ \hline Cℓ_2O (g) & 80.3 \\ \hline F_2O (g) & – 21.7 \\ \hline F_2 (g) & 0 \\ \hline \end{array}
Thermochemistry: Calorimetry
Ethene gas (C2H4) reacts with fluorine gas according to the following balanced equation:
C2H4(g)+6F2(g)→2CF4(g)+4HF(g)ΔrH°=−3012 kJ/mol⋅rxn @ 25°CC_2H_4 (g) + 6 F_2 (g) → 2 CF_4 (g) + 4 HF (g) \qquadΔ_rH° = −3012\ kJ/mol·rxn\ @ \ 25°C
Substance△fH° @ 25°C(kJ/mol)C2H4(g)??F2(g)0CF4(g)–933HF(g)–273.3\def\arraystretch{1.5} \begin{array}{|c|c|} \hline \rm{Substance} & △_fH° \ @ \ 25°C(kJ/mol) \\ \hline C_2H_4 (g) & ?? \\ \hline F_2 (g) & 0 \\ \hline CF_4 (g) & – 933 \\ \hline HF (g) & – 273.3 \\ \hline \end{array}
Thermodynamics: Enthalpy of reaction
Find ΔHo for the following reaction if 12.0 g of N2H4 is consumed in excess N204.
2N2H4(l)+N2O4(g)→3N2(g)+4H2O(l)2N_2H_{4(l)}+N_2O_{4(g)}\to 3N_{2(g)}+4H_2O_{(l)}
The following enthalpies of formation are provided:
Calculate ΔHof of SO2(g) given the following: (answer in kJ)
S(s)+1.5O2(g)→SO3(g),ΔH=−395.8kJS_{(s)}+1.5O_{2(g)}\to SO_{3(g)}, \Delta H=-395.8kJ
2SO2(g)+O2(g)→2SO3(g),ΔH=−198.2kJ2SO_{2(g)}+O_{2(g)}\to 2SO_{3(g)}, \Delta H=-198.2kJ
Intramolecular bonding: Strength of Covalent Bonds
Calculate the enthalpy for the following reaction.
2 C2H6 (g) + 7 O=O (g) → 4 O=C=O (g) + 6 H-O-H (g)
The following are bond dissociation energies (kJ/mol):
Find the standard molar enthalpy of combustion of C2H5OH using the standard enthalpy of formations below.
ΔHf(C2H5OH(l))=−277.7 kJ/mol\Delta H_f(C_2H_5OH_{(l)}) = -277.7\ kJ/mol
ΔHf(CO2(g))=−393.5 kJ/mol\Delta H_f(CO_{2(g)}) = -393.5\ kJ/mol
Thermochemistry Calculations
The standard enthalpy of fusion for water, ΔHofusion = 6.01 kJ/mol. Calculate the energy required (in kJ) to convert 26.0 g of ice at 0oC to liquid water at 20oC.
Heat of Formation
Determine the standard molar enthalpy of formation of solid Al2O3 from the following equations.
Al (s) + 3 H2O (l) → Al(OH)3 (s) + 3/2 H2 (g) ΔH =+100 kJmol\Delta H\ =+100\ \frac{kJ}{mol}
H2 (g) + 1/2 O2 (g) → H2O (l)ΔH=−300 kJmol\Delta H=-300\ \frac{kJ}{mol}