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Find the standard molar enthalpy of combustion of C2H5OH using the standard ent…

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Wize University Chemistry Textbook > Thermodynamics: Part 1: (0th law, 1st law, Calorimetry, Enthalpy)

Heats of Formation Method

4 Activities

Wize High School Grade 12 Chemistry Textbook > Energy Changes

Ways to Calculate Enthalpy of Reactions

10 Activities

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
ΔHf(H2O(g))=−241.8 kJ/mol\Delta H_f(H_2O_{(g)}) = -241.8\ kJ/mol

Report your answer to four significant figure in kJ/mol. Do not include units in the answer field
More Heats of Formation Method Questions:
The standard enthalpy of formation for the amino acid glycine C2H5NO2(s) is –528 kJ/mol at 25°C. Which reaction below has an enthalpy change equal to this value?
Which of the following represents a proper formation reaction for the product indicated? (select all that apply)
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}
Heats of Formation

Heats of Formation
How much heat is required for the following reaction to go given that we start with 37g of CO2(g)?
CO2(g) + 2H2O(l) --> CH4(g) + 2O2(g) ∆H= +890.4kJ/mol
Heats of Formation

Enthalpy of Formation
Calculate ΔHf\Delta H_f of SO2SO_2
S+1.5O2→SO3ΔH=−395.8kJ2SO2+O2→2SO3ΔH=−198.2kJ\begin{array}{}S+1.5 O_2\rightarrow SO_3\quad \Delta H=-395.8 kJ\\2SO_2+O_2\rightarrow 2SO_3\quad \Delta H=-198.2kJ\end{array}
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:
Find the ΔH for the atomization of gaseous propane (CH3CH2CH3). The bond enthalpy for C-H bonds is 414 kJ/mol and the bond enthalpy for C-C bonds is 347 kJ/mol.
Please report your answer to four significant figures in kJ/mol. Do NOT include units in the answer field
Enthalpy of Reactions
Practice: Formation Reaction
Write a reaction for the heat of formation for dimethylaminopyridine, (CH3)2NC5H4N.
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):
Practice: Hydrogenation of Acetylene
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
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}