8. Thermochemistry

Enthalpy of Formation

Enthalpy of Reaction

Acetylene 1C2H21g22 is used for welding because oxyacetylene
is the hottest burning common fuel gas. Using standard
enthalpies of formation, calculate the quantity of heat produced
when 10 g of acetylene is completely combusted in air
under standard conditions.

Acetic acid (CH3CO2H), whose aqueous solutions are known as vinegar, is prepared by reaction of ethyl alcohol (CH3CH2OH) with oxygen: CH3CH2OH(l) + O2(g) → CH3CO2H(l) + H2O(l) Use the following data to calculate ∆H° in kilojoules for the reaction: ∆H°f [CH3CH2OH(l)] = -277.7 kJ/mol ∆H°f [CH3CO2H(l)] = -484.5 kJ/mol ∆H°f [H2O(l)] = -285.8 kJ/mol

Methyl tert-butyl ether (MTBE) is prepared by reaciton of methanol (l) (ΔH°f = -239.2 kJ/mol) with 2-methyl-propene (g), according to the requation  Calculate ΔH°f in kJ/mol for 2-methylpropene

Using values from Appendix C, calculate the value of H for 
each of the following reactions: 
(a) CaO1s2 + 2HF1g2¡ CaF21s2 + H2O1g2

Ethanol 1C2H5OH2 is blended with gasoline as an automobile 
fuel. 
(d) Calculate the mass of CO2 produced 
per kJ of heat emitted.

Use the information in Table 9.2 to calculate ΔH° in kilojoules for the photosynthesis of glucose (C6H12O6) and O2 from CO2 and liquid H2O, a reaciton carried out by all green plants

Calculate ∆H°f in kJ/mol for benzene, C6H6, from the following data: 2 C6H6(l) + 15 O2(g) → 12 CO2(g) + 6 H2O(l) ∆H°=-6534 kJ ∆H°f (CO2) = -393.5 kJ/mol ∆H°f(H2O) = -285.8 kJ/mol

Top fuel dragsters and funny cars burn nitromethane as fuel
according
to the balanced combustion equation:
2 CH3NO2(l ) + 32O2( g)¡2 CO2( g) + 3 H2O(l ) + N2( g)
ΔH °rxn = -1418 kJ
The enthalpy of combustion for nitromethane is -709.2 kJ>mol.
Calculate the standard enthalpy of formation (ΔH °f ) for nitromethane.

Use the data in Appendix B to find standard enthalpies of reaction in kilojoules for the following processes: (b) 2 H2O2 (aq) → 2 H2O (l) + O2(g)

Hydrazine (N2H4) is a fuel used by some spacecraft. It is normally
oxidized by N2O4 according to the equation:
N2H4(l ) + N2O4( g)¡2 N2O( g) + 2 H2O( g)
Calculate ΔH °rxn for this reaction using standard enthalpies of
formation.

C (s, graphite) + 2 H2 (g) + &#189; O2 (g) &#8594; CH3OH (l)

Correct, the reaction gives one mole of the product for the component elements in their standard states.

2 C (s, graphite) + 4 H2 (g) + O2 (g) &#8594; 2 CH3OH (l)

The heat of formation reaction only involves one mole of the product.

C (s, graphite) + H2 (g) + O2 (g) &#8594; CH3OH (l)

C (s, graphite) + 2 H2 (g) + &#189; O2 (l) &#8594; CH3OH (l)

Oxygen is not a liquid in its standard state.

C (s, graphite) + 4 H (g) + &#189; O2 (g) &#8594; CH3OH (l)

Hydrogen is a diatomic molecule in its standard state.

 A formation reaction expresses the formation of one mole of product directly from its elements in their standard states. 

In which of the following is &#916;H&#176;rxn = &#916;H&#176;f at 1 atm and 298 K?

Use the data in Appendix B to find standard enthalpies of reaction in kilojoules for the following processes: (c) Fe2O3 (s) + 3 CO (g) → Fe(s) + 3 CO2(g)

Methanol 1CH3OH2 can be made by the reaction of CO
with H2:
CO1g2 + 2 H21g2 Δ CH3OH1g2
(a) Use thermochemical data in Appendix C to calculate
H for this reaction.

Ethanol (C2H5OH) can be made from the fermentation of crops
and has been used as a fuel additive to gasoline. Write a balanced
equation for the combustion of ethanol and calculate ΔH °rxn.

The Balance Between Enthalpy and Entropy

The standard enthalpies of formation of gaseous propyne 
1C3H42, propylene 1C3H62, and propane 1C3H82 are 
+185.4, +20.4, and -103.8 kJ>mol, respectively. (c) Which is the most 
efficient fuel in terms of heat evolved per unit mass?

Calcualte ΔH° in kilojoules for the synthesis of lime (CaO) from limestone (CaCO3), the key step in the manufacture of cement. CaCO3(s) → CaO(s) + CO2(g) ΔH°f [CaCO3(s)] = -1207.6 kJ/mol ΔH°f [CaO(s)] = -634.9 kJ/mol ΔH°f [CO2(g)] = -393.5 kJ/mol

Styrene (C8H8), the precursor of polystyrene polymers, has a standard heat of combustion of -4395 kJ/mol. Write a balanced equation for the combustion reaction, and calculate ΔH°f for styrene in kJ/mol. ΔH°f [CO2(g)] = -393.5 kJ/mol; ΔH°f [H2O(l)] = -285.8 kJ/mol

Methanol 1CH3OH2 is made industrially in two steps from CO and H2. It is so cheap to make that it is being considered for use as a precursor to hydrocarbon fuels, such as methane 1CH42:
Step 1. CO1g2 + 2 H21g2 S CH3OH1l2
ΔS° = - 332 J>K
Step 2. CH3OH1l2 S CH41g2 + 1>2 O21g2
ΔS° = 162 J>K
(a) Calculate ΔH° in kilojoules for step 1.

The explosive nitroglycerin (C3H5N3O9) decomposes rapidly
upon ignition or sudden impact according to the balanced
equation:
4 C3H5N3O9(l )¡12 CO2( g) + 10 H2O( g) + 6 N2( g) + O2( g)
ΔH °rxn = -5678 kJ
Calculate the standard enthalpy of formation (ΔH °f ) for
nitroglycerin.

Use standard enthalpies of formation to calculate ΔH °rxn for
each reaction.
a. C2H4( g) + H2( g)¡C2H6( g)

Consider the following equilibrium between oxides of
nitrogen
3 NO1g2 Δ NO21g2 + N2O1g2
(a) Use data in Appendix C to calculate H for this reaction.

Hydrogen and methanol have both been proposed as alternatives to hydrocarbon fuels. To compare the energy of these fuels to that of octane (C8H18), calculate the energy released per kilogram of octane.

Acetylene 1C2H22 and nitrogen 1N22 both contain a triple 
bond, but they differ greatly in their chemical properties. 
(b) By referring to Appendix C, look up the enthalpies of formation 
of acetylene and nitrogen. Which compound is more stable?

When magnesium metal is burned in air (Figure 3.6), two 
products are produced. One is magnesium oxide, MgO. The 
other is the product of the reaction of Mg with molecular 
nitrogen, magnesium nitride. When water is added to magnesium 
nitride, it reacts to form magnesium oxide and ammonia 
gas. (e) The 
standard enthalpy of formation of solid magnesium nitride is 
-461.08 kJ>mol. Calculate the standard enthalpy change for 
the reaction between magnesium metal and ammonia gas.

Pentane (C5H12) is a component of gasoline that burns according
to the following balanced equation:
C5H12(l ) + 8 O2( g)¡5 CO2( g) + 6 H2O( g)
Calculate ΔH °rxn for this reaction using standard enthalpies of
formation. (The standard enthalpy of formation of liquid pentane
is -146.8 kJ>mol.)

Use standard enthalpies of formation to calculate the standard
change in enthalpy for the melting of ice. (The ΔH °f for H2O(s)
is -291.8 kJ/mol.)

It is interesting to compare the “fuel value” of a hydrocarbon
in a hypothetical world where oxygen is not the
combustion agent. The enthalpy of formation of CF41g2 is
-679.9 kJ>mol. Which of the following two reactions is the
more exothermic?
CH41g2 + 2 O21g2 ¡ CO21g2 + 2 H2O1g2
CH41g2 + 4 F21g2 ¡ CF41g2 + 4 HF1g2

Use standard enthalpies of formation to calculate ΔH °rxn for 
each reaction. 
d. Cr2O3(s) + 3 CO( g)¡2 Cr(s) + 3 CO2( g)

Under certain nonstandard conditions, oxidation by O2( g) of
1 mol of SO2( g) to SO3( g) absorbs 89.5 kJ. The enthalpy of formation
of SO3( g) is -204.2 kJ under these conditions. Find the
enthalpy of formation of SO2( g).

Ethanol 1C2H5OH2 is blended with gasoline as an automobile
fuel. 
(c) Calculate the heat produced per liter of ethanol by combustion
of ethanol under constant pressure. Ethanol has a
density of 0.789 g>mL.

All the oxides of nitrogen have positive values of ΔGf° at 298 K,
but only one common oxide of nitrogen has a positive ΔSf°.
Identify that oxide of nitrogen without reference to thermodynamic data and explain.

The mechanism for the oxidation of HBr by O2 to form
2 H2O and Br2 is shown in Exercise 14.74. (a) Calculate the
overall standard enthalpy change for the reaction process.

Given the standard heats of formation shown in Appendix B, what is ΔH° in kilojoules for the reaciton 3 N2O4(g) + 2 H2O(l) → 4 HNO3(aq) + 2 NO(g)

Consider the following equation:&nbsp;2 ClF3(g) + 2 NH3(g) → 1 N2(g) + 6 HF (g) + 6 Cl2(g) &nbsp;ΔHrxn = –1196 kJDetermine the standard enthalpy of formation for chlorine trifluoride, ClF3.&nbsp;<img src="https://static.studychannel.pearsonprd.tech/courses/general-chemistry/thumbnails/images/7d4dfb2d"alt=""/>

The fuel in high-efficiency natural-gas vehicles consists primarily of methane 1CH42. (a) How much heat is produced in burning 1 mol of CH41g2 under standard conditions if reactants and products are brought to 298 K and H2O1l2 is formed?

Hydrogen and methanol have both been proposed as alternatives to hydrocarbon fuels. Which fuel contains the most energy in the least mass?

Combustion analysis of 0.1500 g of methyl tert-butyl 
ether, an octane booster used in gasoline, gave 0.3744 g of 
CO2 and 0.1838 g of H2O. When a flask having a volume 
of 1.00 L was evacuated and then filled with methyl tertbutyl 
ether vapor at a pressure of 100.0 kPa and a temperature 
of 54.8 °C, the mass of the flask increased by 3.233 g. 
 
(d) The enthalpy of combustion for methyl tert-butyl ether 
is ΔH° combustion = -3368.7 kJ>mol. What is its standard 
enthalpy of enthalpy of formation, ΔH°f?

Methyl tert-butyl ether (MTBE), C5H12O, a gasoline additive used to boost octane ratings, has ΔH°f = -313.6 kJ/mol. Write a balanced equaiton for its combustion reaciton, and calcualte its standard heat of combustion in kilojoules

Methanol 1CH3OH2 is made industrially in two steps from CO and H2. It is so cheap to make that it is being considered for use as a precursor to hydrocarbon fuels, such as methane 1CH42:
Step 1. CO1g2 + 2 H21g2 S CH3OH1l2
ΔS° = - 332 J>K
Step 2. CH3OH1l2 S CH41g2 + 1>2 O21g2
ΔS° = 162 J>K (f) Calculate ΔH° for step 2.

Use the data in Appendix B to find standard enthalpies of reaction in kilojoules for the following processes: (a) C(s) + CO2(g) → 2 CO(g)

Zinc metal reacts with hydrochloric acid according to the balanced
equation:
Zn(s) + 2 HCl(aq)¡ZnCl2(aq) + H2( g)
When 0.103 g of Zn(s) is combined with enough HCl to make 50.0
mL of solution in a coffee-cup calorimeter, all of the zinc reacts,
raising the temperature of the solution from 22.5 °C to 23.7 °C.
Find ΔHrxn for this reaction as written. (Use 1.0 g>mL for the density
of the solution and 4.18 J>g # °C as the specific heat capacity.)

Use standard enthalpies of formation to calculate ΔH °rxn for 
each reaction. 
c. 3 NO2( g) + H2O(l )¡2 HNO3(aq) + NO( g)

A table of standard enthalpies of formation (ΔH°f) gives a value of −467.9 kJ/mol for NaNO3(s). Which reaction has a ΔH° value of −467.9 kJ?
(a) Na+ (aq) + NO3−(aq) → NaNO3(s)
(b) Na(s) + N(g + O3(g) → NaNO3(s)
(c) Na(s) + 1/2 N2(g) + 3/2 O2(g) → NaNO3(s)
(d) 2 Na(s) + N2(g) + 3 O2(g) → 2 NaNO3(s)

The ΔH for the oxidation of sulfur in the gas phase to SO3 is
-204 kJ/mol and for the oxidation of SO2 to SO3 is 89.5 kJ/mol.
Find the enthalpy of formation of SO2 under these conditions.

Use standard enthalpies of formation to calculate ΔH °rxn for
each reaction.
a. 2 H2S( g) + 3 O2( g)¡2 H2O(l ) + 2 SO2( g)

Isooctane, C8H18, is the component of gasoline from which the term octane rating derives (b) The standard molar heat of combustion of isooctane (l) is -5461 kJ/mol. Calculate ΔH°f for isooctane(l)

Correct application of standard heat of formation values.

In this case, the summation of the reactants' heats of formation is a positive number that is subtracted for that of the products (which is a negative value).

Water in the product is in the liquid state.

All the coefficients in the equation must be considered in the calculation.

 Enthalpy is a state function, therefore the change in the enthalpy of reaction can be calculated by finding the difference in the sum of the change in enthalpy of formation of the products and the sum of the change in enthalpy of formation of the reactants. 

Using standard thermodynamic data, calculate &#916;H&#176;rxn for the following: 2 C2H2 (g) + 5 O2 (g) &#8594; 4 CO2 (g) + 2 H2O (l)

Calcium carbide 1CaC22 reacts with water to form acetylene
1C2H22 and Ca1OH22. From the following enthalpy of reaction
data and data in Appendix C, calculate Hf° for CaC21s2:
CaC21s2 + 2 H2O1l2 ¡ Ca1OH221s2 + C2H21g2
H° = -127.2 kJ

Gasoline is composed primarily of hydrocarbons, including 
many with eight carbon atoms, called octanes. One of 
the cleanest–burning octanes is a compound called 2,3,4- 
trimethylpentane, which has the following structural formula: 
 
The complete combustion of one mole of this compound 
to CO21g2 and H2O1g2 leads to H° = -5064.9 kJ. 
(b) By using the information in 
this problem and data in Table 5.3, calculate Hf° for 
2,3,4-trimethylpentane.

Metallic mercury is obtained by heating the mineral cinnabar (HgS) in air:
HgS1s2 + O21g2 S Hg1l2 + SO21g2
(a) Use the data in Appendix B to calculate ΔH° in kilojoules for the reaction.

Consider the following acid-neutralization reactions involving
the strong base NaOH(aq):
HNO31aq2 + NaOH1aq2¡NaNO31aq2 + H2O1l2
HCl1aq2 + NaOH1aq2¡NaCl1aq2 + H2O1l2
NH4+1aq2 + NaOH1aq2¡NH31aq2 + Na+1aq2 + H2O1l2
(d) In the third equation NH4
+1aq2 is acting as an acid. Based
on the value of H° for this reaction, do you think it is a
strong or a weak acid? Explain.

One of the best-selling light, or low-calorie, beers is 4.2%
alcohol by volume and a 355-mL serving contains 110
Calories; remember: 1 Calorie = 1000 cal = 1 kcal. To
estimate the percentage of Calories that comes from
the alcohol, consider the following questions.
(a) Write
a balanced chemical equation for the reaction of ethanol,
C2H5OH, with oxygen to make carbon dioxide
and water. 
(b) Use enthalpies of formation in Appendix
C to determine ΔH for this reaction. 
(c) If 4.2% of
the total volume is ethanol and the density of ethanol
is 0.789 g/mL, what mass of ethanol does a 355-mL serving
of light beer contain? 
(d) How many Calories are
released by the metabolism of ethanol, the reaction from
part (a)? 
(e) What percentage of the 110 Calories comes
from the ethanol?

The methane molecule, CH4, has the geometry shown in
Figure 2.17. Imagine a hypothetical process in which the
methane molecule is “expanded,” by simultaneously extending
all four C—H bonds to infinity. We then have the
process
CH41g2¡C1g2 + 4 H1g2
(a) Compare this process with the reverse of the reaction
that represents the standard enthalpy of formation of
CH41g2.

Acid spills are often neutralized with sodium carbonate or sodium hydrogen carbonate. For neutralization of acetic acid, the unbalanced equations are
112 CH3CO2H1l2 + Na2CO31s2 S
CH3CO2Na1aq2 + CO21g2 + H2O1l2 122 CH3CO2H1l2 + NaHCO31s2
CH3CO2Na1aq2 + CO21g2 + H2O1l2
(c) How much heat in kilojoules is absorbed or liberated in each reaction? See Appendix B for standard heats of for- mation; ΔH°f = - 726.1 kJ>mol for CH3CO2 Na(aq).

Use standard enthalpies of formation to calculate ΔH °rxn for
each reaction.
b. SO2( g) + 12 O2( g)¡SO3( g)

(b) Use the data in Appendix B to calculate ΔH° for the reaction of potassium metal with water.

Given the standard heats of formation shown in Appendix B, what is ΔH° in kilojoules for the reaciton CaCO3(s) → CaO(s) + CO2(g)?

Methanol 1CH3OH2 is used as a fuel in race cars. 
(b) Calculate the standard enthalpy change for the reaction, 
assuming H2O1g2 as a product.

When a gaseous compound X containing only C, H, and O
is burned in O2, 1 volume of the unknown gas reacts with
3 volumes of O2 to give 2 volumes of CO2 and 3 volumes
of gaseous H2O. Assume all volumes are measured at the
same temperature and pressure.
(d) Combustion of 5.000 g of X releases 144.2 kJ heat.
Look up ΔH°f values for CO21g2 and H2O1g2 in
Appendix B, and calculate ΔH°f for compound X.

Atomic hydrogen (H) is used in welding (AHW). The atoms
recombine to hydrogen molecules with a large release of heat
according to the following reaction:
2 H1g2 ¡H21g2
(b) Which has the higher enthalpy under these conditions,
2 H1g2 or H21g2?

The oxidation of ammonia is illustrated by the following equation:&nbsp;<img src="https://static.studychannel.pearsonprd.tech/courses/general-chemistry/thumbnails/images/c26f007f"class="fr-fic fr-dii">Calculate the enthalpy of reaction,&nbsp;ΔHRxn, based on the given standard heats of formation.&nbsp;<img src="https://cdn.clutchprep.com/problem_images/147428-f7a4076eea0495b8.png"alt=""/>

Complete combustion of 1 mol of acetone 1C3H6O2 liberates
1790 kJ:
C3H6O1l2 + 4 O21g2 ¡ 3 CO21g2 + 3 H2O1l2
H° = -1790 kJ
Using this information together with the standard enthalpies
of formation of O21g2, CO21g2, and H2O1l2 from
Appendix C, calculate the standard enthalpy of formation of
acetone.

Enthalpy Stoichiometry Part 1: Finding Heat and Mass

What is ΔH for the explosion of nitroglycerin? (LO 9.14)
2 C3H5(NO3)3(l) → 3 N2(g) + 1/2 O2(g) + 6 CO2(g) + 5 H2O(g)
<TABLE>
(a) −315.0 kJ
(b) −4517 kJ
(c) −3425 kJ
(d) −3062 kJ

Correct calculation and manipulation of the equations.

One of the reactions has to be reversed and this changes of the sign of &#916;H.

One of the reactions has to be reversed and this changes of the sign of &#916;H and the given reaction contains two moles of NO2.

If a chemical reaction is expressed as the sum of a series of steps, then &#916;Hrxn for the overall equation is the sum of the heats of reaction for each step.

Find &#916;Hrxn for the following reaction: N2 + 2 O2 &#8594; 2 NO2 Based on the following data. 2 NO &#8594; N2 + O2&nbsp;&nbsp;&nbsp;&nbsp;&#916;H = &#x2212;180 kJ 2 NO + O2 &#8594; 2 NO2&nbsp;&nbsp;&nbsp;&nbsp;&#916;H = &#x2212;112 kJ

What is Enthalpy?

What is a compound’s standard heat of formation?

The standard enthalpies of formation of gaseous propyne
1C3H42, propylene 1C3H62, and propane 1C3H82 are
+185.4, +20.4, and -103.8 kJ>mol, respectively. (b) Calculate the heat evolved on
combustion of 1 kg of each substance.

Methanol (CH3OH) has been suggested as a fuel to replace gasoline.
Find ΔH °rxn, and determine the mass of carbon dioxide
emitted per kJ of heat produced. Use the information from the
previous exercise to calculate the same quantity for octane,
C8H18. How does methanol compare to octane with respect to
global warming?

Enthalpy of Formation Example 1

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1. Intro to General Chemistry

2. Atoms & Elements

3. Chemical Reactions

4. BONUS: Lab Techniques and Procedures

5. BONUS: Mathematical Operations and Functions

6. Chemical Quantities & Aqueous Reactions

7. Gases

9. Quantum Mechanics

10. Periodic Properties of the Elements

11. Bonding & Molecular Structure

12. Molecular Shapes & Valence Bond Theory

13. Liquids, Solids & Intermolecular Forces

14. Solutions

15. Chemical Kinetics

16. Chemical Equilibrium

17. Acid and Base Equilibrium

18. Aqueous Equilibrium

19. Chemical Thermodynamics

20. Electrochemistry

21. Nuclear Chemistry

22. Organic Chemistry

23. Chemistry of the Nonmetals

24. Transition Metals and Coordination Compounds

[BOOK CHEM] Clutch Version

Hess's Law

Hess's Law involves the use of partial reactions in the determination of the overall enthalpy of reaction.

Nature of Energy

Thermochemistry is the study of matter and energy associated with chemical reactions or physical changes.&nbsp;

Endothermic & Exothermic Reactions

Endothermic and Exothermic Reactions involve the absorbing and releasing of thermal energy respectively.&nbsp;&nbsp;

Heat Capacity

Heat Capacity is the amount of heat required to change the temperature of a substance.&nbsp;

Thermochemical Equations

Thermochemical Equations&nbsp;involve a balanced chemical equation with a given enthalpy value.

Internal Energy

The Internal Energy of a system is the total energy from all forms of kinetic and potential energy.&nbsp;

Kinetic & Potential Energy

Mechanical Energy is the energy associated with an object's velocity (Kinetic Energy) and its position (Potential Energy).&nbsp;

Formation Equations

Formation Equations involve the standard states of elements combining to form 1 mole of product.

The Enthalpy of Formation for an element is a key component in determining the enthalpy of reaction.&nbsp;

Thermal Equilibrium

Thermal Equilibrium involves two substances that are in physical contact reaching the same final temperature over time.&nbsp;

Constant-Pressure Calorimetry

Constant-Pressure Calorimetry uses a coffee cup calorimeter to determine the heat transfers occurring in a liquid solution.

Constant-Volume Calorimetry

Constant-Volume Calorimetry uses a bomb calorimeter and a combustion reaction to determine its enthalpy of reaction.&nbsp;

First Law of Thermodynamics

The First Law of Thermodynamics states that energy cannot be created nor destroyed, but instead is transferred.&nbsp;