11. Bonding & Molecular Structure

Bond Energy

Bond Length and Bond Energy

If hydrogen were used as a fuel, it could be burned according to
this reaction:
H2( g) + 1 2 O2( g)¡H2O( g)
Use average bond energies to calculate ΔHrxn for this reaction.

Calculate an approximate heat of combustion for ethane (C2H6) in kilojoules by using the bond dissocation energies in Table 9.3. (The strength of the O“O bond is 498 kJ/ mol, and that of a C “ O bond in CO2 is 804 kJ/mol.)

Sulfur tetrafluoride 1SF42 reacts slowly with O2 to form sulfur 
tetrafluoride monoxide 1OSF42 according to the following 
unbalanced reaction: 
SF41g2 + O21g2¡OSF41g2 
The O atom and the four F atoms in OSF4 are bonded to a 
central S atom. (c) Use average bond enthalpies (Table 8.3) to estimate 
the enthalpy of the reaction. Is it endothermic or exothermic?

Ethanol is a possible fuel. Use average bond energies to calculate
ΔHrxn for the combustion of ethanol.
CH3CH2OH( g) + 3 O2( g)¡2 CO2( g) + 3 H2O( g)

State whether each of these statements is true or false. 
(e) The longer the bond, the more energy is 
stored chemical bonds.

Bond Strength and Bond Length

The graph shows how potential energy changes as a function of the distance between two atoms. (LO 7.2)
<img src="https://lightcat-files.s3.amazonaws.com/problem_images/3948f88e8b782fe9-1670278159396.jpg" alt="" />
What is the length of the bond between the two atoms? 
(a) 3.4 angstroms
(b) 3.8 angstroms
(c) 6.0 angstroms
(d) 8.0 angstroms

Consider the reaction 2 H21g2 + O21g2¡2 H2O1l2.
(a) Use the bond enthalpies in Table 5.4 to estimate H for
this reaction, ignoring the fact that water is in the liquid
state.

In the Chemistry and the Environment box on free radicals in this
chapter, we discussed the importance of the hydroxyl radical in
reacting with and eliminating many atmospheric pollutants.
However, the hydroxyl radical does not clean up everything. For
example, chlorofluorocarbons—which destroy stratospheric
ozone—are not attacked by the hydroxyl radical. Consider the
hypothetical reaction by which the hydroxyl radical might
react with a chlorofluorocarbon:
OH( g) + CF2Cl2( g)¡HOF( g) + CFCl2( g)
Use bond energies to explain why this reaction is improbable.
(The C¬F bond energy is 552 kJ>mol.)

(a) The nitrogen atoms in an N2 molecule are held together
by a triple bond; use enthalpies of formation in Appendix
C to estimate the enthalpy of this bond, D(N‚N). 
(b) Consider the reaction between hydrazine and hydrogen to
produce ammonia, N2H41g2 + H21g2¡2 NH31g2. Use enthalpies
of formation and bond enthalpies to estimate the enthalpy
of the nitrogen– nitrogen bond in N2H4. 
(c) Based on 
your answers to parts (a) and (b), would you predict that the 
nitrogen–nitrogen bond in hydrazine is weaker than, similar 
to, or stronger than the bond in N2 ?

The P¬P bond length in white phosphorus is 189 pm. The
Cl¬Cl bond length in Cl2 is 199 pm. (b) What bond length is predicted
for PCl3, using the atomic radii in Figure 7.7?

The F ¬ F bond in F2 is relatively weak because the lone pairs of electrons on one F atom repel the lone pairs on the other F atom; Kp = 7.83 at 1500 K for the reaction F21g2 ∆ 2 F1g2.
(c) Why is the F ¬ F bond in F2 weaker than the Cl ¬ Cl bond in Cl2?

The following plot shows the potential energy of two Cl atoms
as a function of the distance between them. (c) If the Cl2 molecule is compressed under higher and higher pressure, does the Cl–Cl bond
become stronger or weaker?
<img src="https://lightcat-files.s3.amazonaws.com/problem_images/009a6028c539ca84-1660918358862.jpg" alt="" />

If hydrogen were used as a fuel, it could be burned according to
this reaction:
H2( g) + 1 2 O2( g)¡H2O( g)
Which fuel yields more energy per gram?

When individual bond energies are given, we will use the first formula.

Consider the following equation:<img src="https://static.studychannel-prd.pearsonprd.tech/courses/general-chemistry/thumbnails/images/781b2bae"class="fr-fic fr-dii">Determine the bond enthalpy value for the F–S bond.<img src="https://lightcat-files.s3.amazonaws.com/problem_images/fd1a8cebd6bb819e-1608583566374.jpg"class="fr-fic fr-dii">

Covalent Bonding: Bond Order, Bond Length, Bond Energy

Two compounds are isomers if they have the same chemical
formula but different arrangements of atoms. Use Table 8.3
to estimate H for each of the following gas-phase isomerization
reactions and indicate which isomer has the lower
enthalpy.
(d) Methyl isocyanide Acetonitrile<img src="https://lightcat-files.s3.amazonaws.com/problem_images/c8083f19baf5b074-1660910794894.jpg" alt="" />

What is the best prediction for the carbon–oxygen bond length in the carbonate anion, CO3 2-? (LO 7.14)
Data for average carbon-oxygen bond lengths: 

(a) 143 pm 
(b) 132 pm
(c) 121 pm
(d) 118 pm

Use average bond enthalpies from Table 8.4 to estimate 
the enthalpies of the following gas-phase reactions: 
Reaction 1: HF1g2 + H2O1g2 Δ F-1g2 + H3O+1g2 
Reaction 2: HCl1g2 + H2O1g2 Δ Cl-1g2 + H3O+1g2 
Are both reactions exothermic? How do these values relate to 
the different strengths of hydrofluoric and hydrochloric acid?

Explain the difference in the bond dissociation energies for the following bonds: (C-F, 450 kJ/mol), (N-F, 270 kJ/mol), (O-F, 180 kJ/mol), (F-F, 159 kJ/mol).

The P¬P bond length in white phosphorus is 189 pm. The
Cl¬Cl bond length in Cl2 is 199 pm. (a) Based on these
data, what is the predicted P¬Cl bond length in phosphorus
trichloride, PCl3, in which each of the three Cl atoms is
bonded to the P atom?

Bond Energy Example 1

Bond Energy & Bond Length, Forces of Attraction & Repulsion - Chemistry

If hydrogen were used as a fuel, it could be burned according to
this reaction:
H2( g) + 1 2 O2( g)¡H2O( g)
Which fuel yields more energy per mole?

Calculate ΔHrxn for the combustion of octane (C8H18), a component
of gasoline, by using average bond energies and then calculate
it using enthalpies of formation from Appendix IIB.
What is the percent difference between your results? Which result would you expect to be more accurate?

Hydrogenation reactions are used to add hydrogen across double
bonds in hydrocarbons and other organic compounds. Use
average bond energies to calculate ΔHrxn for the hydrogenation
reaction.
H2C“CH2( g) + H2( g)¡H3C¬CH3( g)

Phosgene, COCl21g2, is a toxic gas used as an agent of warfare in World War I.
(b) Using the table of bond dissociation energies (Table 9.3) and the value ΔH°f = 716.7 kJ>mol for C(g), estimate ΔH° f for COCl21g2 at 25 °C. Compare your answer to the actual ΔH° f given in Appendix B, and explain why your calculation is only an estimate.

Bond Strength and Length

Use bond enthalpies in Table 5.4 to estimate H for each of 
the following reactions:
(a)

(a) In polyvinyl chloride shown in Table 12.6, which 
bonds have the lowest average bond enthalpy?<img src="https://lightcat-files.s3.amazonaws.com/problem_images/f3ba5343c92d4a4f-1665510042561.jpg" alt="" />

If hydrogen were used as a fuel, it could be burned according to
this reaction:
H2( g) + 1 2 O2( g)¡H2O( g)
Use average bond energies to calculate ΔHrxn for the combustion of methane (CH4).

The reaction S81g2 S 4 S21g2 has ΔH° = + 237 kJ
(b) The average S ¬ S bond dissociation energy is 225 kJ/mol. Using the value of ΔH° given above, what is the S “ S double bond energy in S21g2?

Calculate the heat of atomization (see previous problem) of 
C2H3Cl, using the average bond energies in Table 9.3.

Bond Energies

(a) Compare the bond enthalpies (Table 8.3) of the carbon–
carbon single, double, and triple bonds to deduce an average
p@bond contribution to the enthalpy. What fraction of
a single bond does this quantity represent?

(b) Based on aver- age bond enthalpies, would you expect a photon capable of dissociating a C ¬ Cl bond to have sufficient energy to dis- sociate a C ¬ Br bond?

Use the data in Table 9.3 to calculate an approximate ∆H° in kilojoules for the synthesis of hydrazine from ammonia: 2 NH3(g) + Cl2(g) → N2H4(g) + 2 HCl(g)

(b) When subjected to high pressure and heated, polyvinyl
chloride converts to diamond. During this transformation
which bonds are most likely to break first?<img src="https://lightcat-files.s3.amazonaws.com/problem_images/94ac455f1c2f6a07-1665510294516.jpg" alt="" />

The heat of atomization is the heat required to convert a molecule
in the gas phase into its constituent atoms in the gas phase.
The heat of atomization is used to calculate average bond energies.
Without using any tabulated bond energies, calculate the
average C¬Cl bond energy from the following data: the heat
of atomization of CH4 is 1660 kJ>mol, and the heat of atomization
of CH2Cl2 is 1495 kJ>mol.

Use the bond energies to estimate the enthalpy of reaction for the combustion of 5 moles of acetylene:<img src="https://static.studychannel-prd.pearsonprd.tech/courses/general-chemistry/thumbnails/images/e38da23f"class="fr-fic fr-dii">

(a) Use average bond enthalpies (Table 8.3) to estimate H 
for the atomization of naphthalene, C10H8:

Isomers are molecules that have the same chemical formula but different arrangements of atoms, as shown here for two isomers of pentane, C5H12.
<img src="https://lightcat-files.s3.amazonaws.com/problem_images/e2a35b4574541839-1665501700151.jpg" alt="" />
(a) Do you expect a significant difference in the enthalpy of combustion of the two isomers? Explain.

Consider the collection of nonmetallic elements: B, As, O, 
and I. (b) Which two would form the longest single bond?

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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

8. Thermochemistry

9. Quantum Mechanics

10. Periodic Properties of the Elements

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

Lewis Dot Structures: Neutral Compounds

Lewis Dot Structures or Electron Dot Structures&nbsp;are diagrams that show how elements in a molecule use their valence electrons to form bonds.&nbsp;

Average Bond Order

Average Bond Order represents the average number of chemical bonds between a pair of bond elements.

Lewis Dot Structures: Ions

Lewis Dot Structures of Ions involves losing or gaining valence electrons to draw the most likely structure.&nbsp;

Octet Rule

The&nbsp;Octet Rule states that main-group elements will generally form enough bonds to obtain 8 electrons in their valence shell.&nbsp;

Lewis Dot Symbols

Lewis Dot Symbols represent the outer shell electrons of an element.&nbsp;

Lewis Dot Structures: Sigma & Pi Bonds

Sigma Bonds directly connect element together, while Pi Bonds provide insulation and protection for those bonds.&nbsp;

Lewis Dot Structures: Exceptions

Lewis Dot Structure Exceptions&nbsp;happen when the central element violates the Octet Rule.

Born Haber Cycle

In the Born-Haber Cycle, ionic solids are created through the ionization and eventual combination of elements.&nbsp;

Coulomb's Law

Coulomb's Law is used to determine the forces experienced by charged particles.&nbsp;

Chemical Bonds

A chemical bond is the connection&nbsp;of different elements through the sharing, transferring or pooling of electrons.&nbsp;

Resonance Structures

Resonance Structures are used to represent the different possible bonding arrangements of a molecule.

Lewis Dot Structures: Acids

Drawing the Lewis Dot Structure of Acids&nbsp;involves first drawing the polyatomic anion and then attaching the H+ ion.&nbsp;

Formal Charge

Formal Charge is the charge given to an element within a molecule when assuming equal sharing of electrons between bonded elements.&nbsp;

Bond Energy is the amount of energy stored within a chemical bond.

Lattice Energy

Lattice Energy&nbsp;is the estimate for the strength of ionic bonds.&nbsp;

Dipole Moment

Dipole arrows are used anytime a molecule possesses a dipole moment, which happens when a molecule is polar.&nbsp;