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Chemistry: The Central Science

Theodore E. Brown, Eugene H LeMay, Bruce E. Bursten

Chapter 8

Basic Concepts of Chemical Bonding - all with Video Answers

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Chapter Questions

01:23

Problem 1

For each of these Lewis symbols, indicate the group in the periodic table in which the element $X$ belongs:
(a) $\cdot \dot{\mathrm{X}} \cdot(\mathrm{b}) \cdot \mathrm{X} \cdot(\mathrm{c}): \dot{\mathrm{X}}$

Sima Sarker
Sima Sarker
Numerade Educator
01:13

Problem 2

Illustrated at right are four ions $-\mathrm{A}, \mathrm{B}, \mathrm{X}$, and $\mathrm{Y}-$ showing their relative ionic radii. The ions shown in red carry positive charges: a $2+$ charge for $A$ and a $1+$ charge for $\mathrm{B}$. Ions shown in blue carry negative charges:
a $1-$ charge for $X$ and a 2 - charge for $Y$. (a) Which combinations of these ions produce ionic compounds where there is a $1: 1$ ratio of cations and anions? (b) Among those compounds, which combination of ions leads to the ionic compound having the largest lattice energy?
(c) Which combination of ions leads to the ionic compound having the smallest lattice energy? [Section 8.2]

Lottie Adams
Lottie Adams
Numerade Educator
02:16

Problem 3

The orbital diagram below shows the valence electrons for a $2+$ ion of an element. (a) What is the element?
(b) What is the electron configuration of an atom of this element? [Section 8.2]

Sima Sarker
Sima Sarker
Numerade Educator
01:53

Problem 4

In the Lewis structure shown below, $A, D, E, Q, X$, and $Z$ represent elements in the first two rows of the periodic table (H-Ne). Identify all six elements so that the formal charges of all atoms are zero. [Section 8.3]

Susan Hallstrom
Susan Hallstrom
Numerade Educator
01:12

Problem 5

The partial Lewis structure below is for a hydrocarbon molecule. In the full Lewis structure, each carbon atom satisfies the octet rule, and there are no unshared electron pairs in the molecule. The carbon-carbon bonds are labeled 1, 2, and 3. (a) Determine where the hydrogen atoms are in the molecule. (b) Rank the carbon-carbon bonds in order of increasing bond length. (c) Rank the carbon-carbon bonds in order of increasing bond enthalpy. [Section $8.3$ and $8.81$

Lottie Adams
Lottie Adams
Numerade Educator
12:09

Problem 6

Consider the Lewis structure for the polyatomic oxyanion shown below, where $X$ is an element from the 3rd period (Na- Ar). By changing the overall charge, $n$, from $1-$ to $2-$ to $3-$ we get three different polyatomic ions. For each of these ions (a) Identify the central atom, $\mathrm{X}$.
(b) Determine the formal charge of the central atom, $\mathrm{X}$. (c) Draw a Lewis structure that makes the formal charge on the central atom equal to zero. (d) If the Lewis structure you drew in part (c) differs from the Lewis structure shown below, which one do you think is the best one (if you were able to choose only a single Lewis structure)? [Sections 8.5, 8.6, and 8.7]

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:24

Problem 7

(a) What are valence electrons? (b) How many valence electrons does a nitrogen atom possess? (c) An atom has the electron configuration $1 s^{2} 2 s^{2} 2 p^{6} 3 s^{2} 3 p^{2} .$ How many valence electrons does the atom have?

Narayan Hari
Narayan Hari
Numerade Educator
01:49

Problem 8

(a) What is the octet rule? (b) How many electrons must a sulfur atom gain to achieve an octet in its valence shell? (c) If an atom has the electron configuration $1 s^{2} 2 s^{2} 2 p^{3}$, how many electrons must it gain to achieve an octet?

Nicole Smina
Nicole Smina
Numerade Educator
02:40

Problem 9

Write the electron configuration for phosphorus. Identify the valence electrons in this configuration and the nonvalence electrons. From the standpoint of chemical reactivity, what is the important difference between them?

Sima Sarker
Sima Sarker
Numerade Educator
03:36

Problem 10

(a) Write the electron configuration for the element titanium, Ti. How many valence electrons does this atom possess?
(b) Hafnium, Hf, is also found in group $4 \mathrm{~B}$. Write the electron configuration for Hf. (c) Both Ti and Hf behave as though they possess the same number of valence electrons. Which of the subshells in the electron configuration of Hf behave as valence orbitals? Which behave as core orbitals?

James Irizarry
James Irizarry
Numerade Educator
00:50

Problem 11

Write the Lewis symbol for atoms of each of the following elements: (a) $\mathrm{Al}$, (b) $\mathrm{Br}$, (c) $\mathrm{Ar}$, (d) $\mathrm{Sr}$.

Kevin Chimex
Kevin Chimex
Numerade Educator
04:48

Problem 12

What is the Lewis symbol for each of the following atoms or ions: (a) $\mathrm{Ca}$, (b) $\mathrm{P}_{,}(\mathrm{c}) \mathrm{Mg}^{2+}$, (d) $\mathrm{S}^{2-}$ ?

Ishu Khandelwal
Ishu Khandelwal
Numerade Educator
01:01

Problem 13

Using Lewis symbols, diagram the reaction between magnesium and oxygen atoms to give the ionic substance $\mathrm{MgO}$.

Narayan Hari
Narayan Hari
Numerade Educator
01:01

Problem 14

Use Lewis symbols to represent the reaction that occurs between $\mathrm{Ca}$ and $\mathrm{F}$ atoms.

Narayan Hari
Narayan Hari
Numerade Educator
01:41

Problem 15

Predict the chemical formula of the ionic compound formed between the following pairs of elements: (a) $\mathrm{Al}$ and $\mathrm{F}$, (b) $\mathrm{K}$ and $\mathrm{S}$, (c) $\mathrm{Y}$ and $\mathrm{O}$, (d) $\mathrm{Mg}$ and $\mathrm{N}$.

Kevin Chimex
Kevin Chimex
Numerade Educator
01:18

Problem 16

Which ionic compound is expected to form from combining the following pairs of elements: (a) barium and fluorine, (b) cesium and chlorine, (c) lithium and nitrogen, (d) aluminum and oxygen?

James Irizarry
James Irizarry
Numerade Educator
05:33

Problem 17

Write the electron configuration for each of the following ions, and determine which ones possess noble-gas configurations:
(a) $\mathrm{Sr}^{2+}$,
(b) $\mathrm{Ti}^{2+}$, (c) $\mathrm{Se}^{2-}$,
(d) $\mathrm{Ni}^{2+}$,
(e) $\mathrm{Br}^{-}$, (f) $\mathrm{Mn}^{3+}$.

Kevin Chimex
Kevin Chimex
Numerade Educator
03:52

Problem 18

Write electron configurations for the following ions, and determine which have noble-gas configurations:
(a) $\mathrm{Zn}^{2+}$,
(b) $\mathrm{Te}^{2-}$
(c) $\mathrm{Sc}^{3+}$,
(d) $\mathrm{Rh}^{3+}$,
(e) $\mathrm{Tl}^{+}$, (f) $\mathrm{Bi}^{3+}$.

Sima Sarker
Sima Sarker
Numerade Educator
01:41

Problem 19

(a) Define the term lattice energy. (b) Which factors govern the magnitude of the lattice energy of an ionic compound?

Narayan Hari
Narayan Hari
Numerade Educator
01:46

Problem 20

$\mathrm{NaCl}$ and KF have the same crystal structure. The only difference between the two is the distance that separates cations and anions. (a) The lattice energies of $\mathrm{NaCl}$ and $\mathrm{KF}$ are given in Table $8.2 .$ Based on the lattice energies, would you expect the $\mathrm{Na}-\mathrm{Cl}$ or the $\mathrm{K}-\mathrm{F}$ distance to be longer? (b) Use the ionic radii given in Figure $7.8$ to estimate the $\mathrm{Na}-\mathrm{Cl}$ and K-F distances. Does this estimate agree with the prediction you made based upon the lattice energies?

Lottie Adams
Lottie Adams
Numerade Educator
02:13

Problem 21

The ionic substances $\mathrm{KF}, \mathrm{CaO}$, and $\mathrm{Sc} \mathrm{N}$ are isoelectronic (they have the same number of electrons). Examine the lattice energies for these substances in Table $8.2$, and account for the trends you observe.

Narayan Hari
Narayan Hari
Numerade Educator
02:18

Problem 22

(a) Does the lattice energy of an ionic solid increase or decrease (i) as the charges of the ions increase, (ii) as the sizes of the ions increase? (b) Using a periodic table, arrange the following substances according to their expected lattice energies, listing them from lowest lattice energy to the highest: $\mathrm{ScN}, \mathrm{KBr}, \mathrm{MgO}, \mathrm{NaF}$. Compare your list with the data in Table $8.2$.

Catherine Lemar
Catherine Lemar
Numerade Educator
01:02

Problem 23

The lattice energies of $\mathrm{KBr}$ and $\mathrm{CsCl}$ are nearly equal (Table 8.2). What can you conclude from this observation?

Narayan Hari
Narayan Hari
Numerade Educator
02:02

Problem 24

Explain the following trends in lattice energy: (a) $\mathrm{CaF}_{2}>\mathrm{BaF}_{2} ;$ (b) $\mathrm{NaCl}>\mathrm{RbBr}>\mathrm{CsBr} ;$ (c) $\mathrm{BaO}>\mathrm{KF}$

Catherine Lemar
Catherine Lemar
Numerade Educator
01:28

Problem 25

Energy is required to remove two electrons from Ca to form $\mathrm{Ca}^{2+}$ and is required to add two electrons to $\mathrm{O}$ to form $\mathrm{O}^{2-}$. Why, then, is $\mathrm{CaO}$ stable relative to the free elements?

Narayan Hari
Narayan Hari
Numerade Educator
08:06

Problem 26

Construct a Born-Haber cycle for the formation of the hypothetical compound $\mathrm{NaCl}_{2}$, where the sodium ion has a 2+ charge (the 2nd ionization energy for sodium is given in Table 7.2). (a) How large would the lattice energy need to be for the formation of $\mathrm{NaCl}_{2}$ to be exothermic? (b) If we were to estimate the lattice energy of $\mathrm{NaCl}_{2}$ to be roughly equal to that of $\mathrm{MgCl}_{2}$ ( $2326 \mathrm{~kJ} / \mathrm{mol}$
from Table 8.2), what value would you obtain for the standard enthalpy of formation, $\Delta H_{f}^{\circ}$, of $\mathrm{NaCl}_{2}$ ?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
04:33

Problem 27

Use data from Appendix $C$, Figure $7.12$, and Figure $7.14$ to calculate the lattice energy of RbCl. Is this value greater than or less than the lattice energy of $\mathrm{NaCl}$ ? Explain.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
09:28

Problem 28

(a) Based on the lattice energies of $\mathrm{MgCl}_{2}$ and $\mathrm{SrCl}_{2}$ given in Table $8.2$, what is the range of values that you would expect for the lattice energy of $\mathrm{CaCl}_{2} ?$ (b) Using data from Appendix $C$, Figure $7.12$, and Figure $7.14$ and the value of the second ionization energy for $\mathrm{Ca}$, $1145 \mathrm{~kJ} / \mathrm{mol}$, calculate the lattice energy of $\mathrm{CaCl}_{2} .$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
01:04

Problem 29

(a) What is meant by the term covalent bond? (b) Give three examples of covalent bonding. (c) A substance XY, formed from two different elements, boils at $-33^{\circ} \mathrm{C}$. Is XY likely to be a covalent or an ionic substance? Explain.

Nicole Smina
Nicole Smina
Numerade Educator
01:15

Problem 30

Which of these elements is unlikely to form covalent bonds: $\mathrm{S}, \mathrm{H}, \mathrm{K}, \mathrm{Ar}, \mathrm{Si}$ ? Explain your choices.

Nicole Smina
Nicole Smina
Numerade Educator
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Problem 31

Using Lewis symbols and Lewis structures, diagram the formation of $\mathrm{SiCl}_{4}$ from $\mathrm{Si}$ and $\mathrm{Cl}$ atoms.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:03

Problem 32

Use Lewissymbols and Lewis structures to diagram the formation of $\mathrm{PF}_{3}$ from $\mathrm{P}$ and $\mathrm{F}$ atoms.

Sima Sarker
Sima Sarker
Numerade Educator
04:19

Problem 33

(a) Construct a Lewis structure for $\mathrm{O}_{2}$ in which each atom achieves an octet of electrons. (b) Explain why it is necessary to form a double bond in the Lewis structure. (c) The bond in $\mathrm{O}_{2}$ is shorter than the $\mathrm{O}-\mathrm{O}$ bond in compounds that contain an $\mathrm{O}-\mathrm{O}$ single bond. Explain this observation.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:04

Problem 34

(a) Construct a Lewis structure for hydrogen peroxide, $\mathrm{H}_{2} \mathrm{O}_{2}$, in which each atom achieves an octet of electrons.
(b) Do you expect the $\mathrm{O}-\mathrm{O}$ bond in $\mathrm{H}_{2} \mathrm{O}_{2}$ to be longer or shorter than the $\mathrm{O}-\mathrm{O}$ bond in $\mathrm{O}_{2}$ ?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
01:18

Problem 35

(a) What is meant by the term electronegativity? (b) On the Pauling scale what is the range of electronegativity values for the elements? (c) Which element has the greatest electronegativity? (d) Which element has the smallest electronegativity?

Nicole Smina
Nicole Smina
Numerade Educator
01:39

Problem 36

(a) What is the trend in electronegativity going from left to right in a row of the periodic table? (b) How do electronegativity values generally vary going down a column in the periodic table? (c) How do periodic trends in electronegativity relate to those for ionization energy and electron affinity?

Nicole Smina
Nicole Smina
Numerade Educator
02:48

Problem 37

Using only the periodic table as your guide, select the most electronegative atom in each of the following sets:
(a) Se, $\mathrm{Rb}, \mathrm{O}, \mathrm{In} ;$ (b) $\mathrm{Al}, \mathrm{Ca}, \mathrm{C}, \mathrm{Si} ;$ (c) $\mathrm{Ge}, \mathrm{As}, \mathrm{P}, \mathrm{Sn} ;$ (d) $\mathrm{Li}$, $\mathrm{Rb}, \mathrm{Be}, \mathrm{Sr}$

Catherine Lemar
Catherine Lemar
Numerade Educator
05:03

Problem 38

By referring only to the periodic table, select (a) the most electronegative element in group $6 \mathrm{~A} ;$ (b) the least electronegative element in the group $\mathrm{Al}, \mathrm{Si}, \mathrm{P} ;$ (c) the most electronegative element in the group Ga, $\mathrm{P}, \mathrm{Cl}, \mathrm{Na}$;
(d) the element in the group $\mathrm{K}, \mathrm{C}, \mathrm{Zn}, \mathrm{F}$, that is most likely to form an ionic compound with Ba.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
01:15

Problem 39

Which of the following bonds are polar: (a) $\mathrm{B}-\mathrm{F}$, (b) $\mathrm{Cl}-\mathrm{Cl}$, (c) $\mathrm{Se}-\mathrm{O}$, (d) $\mathrm{H}-\mathrm{I}$ ? Which is the more electronegative atom in each polar bond?

Kevin Chimex
Kevin Chimex
Numerade Educator
04:12

Problem 40

Arrange the bonds in each of the following sets in order of increasing polarity:
(a) $\mathrm{C}-\mathrm{F}, \mathrm{O}-\mathrm{F}, \mathrm{Be}-\mathrm{F}$
(b) $\mathrm{O}-\mathrm{Cl}, \mathrm{S}-\mathrm{Br}, \mathrm{C}-\mathrm{P}$
(c) $\mathrm{C}-\mathrm{S}, \mathrm{B}-\mathrm{F}, \mathrm{N}-\mathrm{O}$

James Irizarry
James Irizarry
Numerade Educator
07:07

Problem 41

The dipole moment and bond distance measured for the highly reactive gas phase OH molecule are $1.78 \mathrm{D}$ and $0.98 \AA$, respectively. (a) Given these values calculate the effective charges on the $\mathrm{H}$ and $\mathrm{O}$ atoms of the OH molecule in units of the electronic charges $e$. (b) Is this bond more or less polar than the $\mathrm{H}-\mathrm{Cl}$ bond in an $\mathrm{HCl}$ molecule? (c) Is that what you would have expected based on electronegativities?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:04

Problem 42

The iodine monobromide molecule, IBr, has a bond length of $2.49 \AA$ and a dipole moment of $1.21 \mathrm{D}$.
(a) Which atom of the molecule is expected to have a negative charge? Explain. (b) Calculate the effective charges on the I and Br atoms in IBr, in units of the electronic charge $e$.

James Irizarry
James Irizarry
Numerade Educator
02:26

Problem 43

In the following pairs of binary compounds determine which one is a molecular substance and which one is an ionic substance. Use the appropriate naming convention (for ionic or molecular substances) to assign a name to each compound: (a) $\mathrm{SiF}_{4}$ and $\mathrm{LaF}_{3}$, (b) $\mathrm{FeCl}_{2}$ and $\operatorname{ReCl}_{6}$, (c) $\mathrm{PbCl}_{4}$ and $\mathrm{RbCl}$.

Kevin Chimex
Kevin Chimex
Numerade Educator
03:02

Problem 44

In the following pairs of binary compounds determine which one is a molecular substance and which one is an ionic substance. Use the appropriate naming convention (for ionic or molecular substances) to assign a name to each compound:
(a) $\mathrm{TiCl}_{4}$ and $\mathrm{CaF}_{2}$,
(b) $\mathrm{ClF}_{3}$ and $\mathrm{VF}_{3}$ '
(c) $\mathrm{SbCl}_{5}$ and $\mathrm{AlF}_{3}$.

James Irizarry
James Irizarry
Numerade Educator
06:44

Problem 45

Draw Lewis structures for the following: (a) $\mathrm{SiH}_{4}$,
(b) $\mathrm{CO}$, (c) $\mathrm{SF}_{2}$,
(d) $\mathrm{H}_{2} \mathrm{SO}_{4}(\mathrm{H}$ is bonded to $\mathrm{O})$,
(e) $\mathrm{ClO}_{2}^{-}$
(f) $\mathrm{NH}_{2} \mathrm{OH}$.

Kevin Chimex
Kevin Chimex
Numerade Educator
11:50

Problem 46

Write Lewis structures for the following: (a) $\mathrm{H}_{2} \mathrm{CO}$ (both $\mathrm{H}$ atoms are bonded to $\mathrm{C}$ ), (b) $\mathrm{H}_{2} \mathrm{O}_{2}$, (c) $\mathrm{C}_{2} \mathrm{~F}_{6}$ (contains a $\mathrm{C}-\mathrm{C}$ bond $),$ (d) $\mathrm{AsO}_{3}{ }^{3-}$, (e) $\mathrm{H}_{2} \mathrm{SO}_{3}(\mathrm{H}$ is bonded to $\mathrm{O})$, (f) $\mathrm{C}_{2} \mathrm{H}_{2}$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
04:37

Problem 47

(a) When talking about atoms in a Lewis structure, what is meant by the term formal charge? (b) Does the formal charge of an atom represent the actual charge on that atom? Explain. (c) How does the formal charge of an atom in a Lewis structure differ from the oxidation number of the atom?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
06:51

Problem 48

(a) Write a Lewis structure for the phosphorus trifluoride molecule, $\mathrm{PF}_{3}$. Is the octet rule satisfied for all the atoms in your structure? (b) Determine the oxidation numbers of the $\mathrm{P}$ and $\mathrm{F}$ atoms. (c) Determine the formal charges of the $\mathrm{P}$ and $\mathrm{F}$ atoms. (d) Is the oxidation number for the $\mathrm{P}$ atom the same as its formal charge? Explain why or why not.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
14:21

Problem 49

Write Lewis structures that obey the octet rule for each of the following, and assign oxidation numbers and formal charges to each atom: (a) $\mathrm{NO}^{+}$, (b) $\mathrm{POCl}_{3}$ (P is bonded to the three $\mathrm{Cl}$ atoms and to the $\mathrm{O}$, (c) $\mathrm{ClO}_{4}^{-}$,
(d) $\mathrm{HClO}_{3}(\mathrm{H}$ is bonded to $\mathrm{O})$.

Heath Mclean
Heath Mclean
Numerade Educator
17:23

Problem 50

For each of the following molecules or ions of sulfur and oxygen, write a single Lewis structure that obeys the octet rule, and calculate the oxidation numbers and formal charges on all the atoms: (a) $\mathrm{SO}_{2}$, (b) $\mathrm{SO}_{3}$, (c) $\mathrm{SO}_{3}{ }^{2-}$. (d) Arrange these molecules/ions in order of increasing S - O bond distance.

Oluwapelumi Kolawole
Oluwapelumi Kolawole
Numerade Educator
05:04

Problem 51

(a) Write one or more appropriate Lewis structures for the nitrite ion, $\mathrm{NO}_{2}^{-} .$ (b) With what allotrope of oxygen is it isoelectronic? (c) What would you predict for the lengths of the bonds in $\mathrm{NO}_{2}^{-}$ relative to $\mathrm{N}-\mathrm{O}$ single bonds?

Sima Sarker
Sima Sarker
Numerade Educator
03:45

Problem 52

Consider the nitryl cation, $\mathrm{NO}_{2}^{+} .$ (a) Write one or more appropriate Lewis structures for this ion. (b) Are resonance structures needed to describe the structure? (c) With what familiar molecule is it isoelectronic?

Heath Mclean
Heath Mclean
Numerade Educator
01:31

Problem 53

Predict the ordering of the $\mathrm{C}-\mathrm{O}$ bond lengths in $\mathrm{CO}$, $\mathrm{CO}_{2}$, and $\mathrm{CO}_{3}^{2-}$

Nicole Smina
Nicole Smina
Numerade Educator
04:31

Problem 54

Based on Lewis structures, predict the ordering of $\mathrm{N}-\mathrm{O}$ bond lengths in $\mathrm{NO}^{+}, \mathrm{NO}_{2}^{-}$, and $\mathrm{NO}_{3}^{-} .$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
05:31

Problem 55

(a) Use the concept of resonance to explain why all six $\mathrm{C}-\mathrm{C}$ bonds in benzene are equal in length. (b) Are the $\mathrm{C}-\mathrm{C}$ bond lengths in benzene shorter than $\mathrm{C}-\mathrm{C} \mathrm{sin}-$ gle bonds? Are they shorter than $\mathrm{C}=\mathrm{C}$ double bonds?

Heath Mclean
Heath Mclean
Numerade Educator
04:20

Problem 56

Mothballs are composed of naphthalene, $\mathrm{C}_{10} \mathrm{H}_{8}$, a molecule of which consists of two six-membered rings of carbon fused along an edge, as shown in this incomplete Lewis structure:
(a) Write two complete Lewis structures for naphthalene. (b) The observed $C-C$ bond lengths in the molecule are intermediate between $\mathrm{C}-\mathrm{C}$ single and $\mathrm{C}=\mathrm{C}$ double bonds. Explain. (c) Represent the resonance in naphthalene in a way analogous to that used to represent it in benzene.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:52

Problem 57

(a) State the octet rule. (b) Does the octet rule apply to ionic as well as to covalent compounds? Explain, using examples as appropriate.

Nicole Smina
Nicole Smina
Numerade Educator
02:13

Problem 58

Considering the nonmetals, what is the relationship between the group number for an element (carbon, for example, belongs to group $4 \mathrm{~A}$; see the periodic table on the inside front cover) and the number of single covalent bonds that element needs to form to conform to the octet rule?

Nicole Smina
Nicole Smina
Numerade Educator
01:54

Problem 59

What is the most common exception to the octet rule? Give two examples.

Sima Sarker
Sima Sarker
Numerade Educator
01:30

Problem 60

For elements in the third row of the periodic table and beyond, the octet rule is often not obeyed. What factors are usually cited to explain this fact?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
07:41

Problem 61

Draw the Lewis structures for each of the following ions or molecules. Identify those that do not obey the octet rule, and explain why they do not. (a) $\mathrm{SO}_{3}{ }^{2-}$, (b) $\mathrm{AlH}_{3}$, (c) $\mathrm{N}_{3}^{-}$, (d) $\mathrm{CH}_{2} \mathrm{Cl}_{2}$, (e) $\mathrm{SbF}_{5}$.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:42

Problem 62

Draw the Lewis structures for each of the following molecules or ions. Which do not obey the octet rule?
(b) $\mathrm{SCN}^{-}$,
(a) $\mathrm{NH}_{4}^{+}$,
(c) $\mathrm{PCl}_{3}$,
(d) $\mathrm{TeF}_{4}$, (e) $\mathrm{XeF}_{2}$.

Crystal Wang
Crystal Wang
Numerade Educator
06:22

Problem 63

In the vapor phase, $\mathrm{BeCl}_{2}$ exists as a discrete molecule. (a) Draw the Lewis structure of this molecule, using only single bonds. Does this Lewis structure satisfy the octet rule? (b) What other resonance forms are possible that satisfy the octet rule? (c) Using formal charges, select the resonance form from among all the Lewis structures that is most important in describing $\mathrm{BeCl}_{2}$.

Sima Sarker
Sima Sarker
Numerade Educator
03:49

Problem 64

(a) Describe the molecule chlorine dioxide, $\mathrm{ClO}_{2}$, using three possible resonance structures. (b) Do any of these resonance structures satisfy the octet rule for every atom in the molecule? Why or why not? (c) Using formal charges, select the resonance structure(s) that is (are) most important.

Crystal Wang
Crystal Wang
Numerade Educator
07:57

Problem 65

Using the bond enthalpies tabulated in Table $8.4$, estimate $\Delta H$ for each of the following gas-phase reactions:

Heath Mclean
Heath Mclean
Numerade Educator
10:35

Problem 66

Using bond enthalpies (Table 8.4), estimate $\Delta H$ for the following gas-phase reactions:

Susan Hallstrom
Susan Hallstrom
Numerade Educator
06:53

Problem 67

Using bond enthalpies (Table $8.4$ ), estimate $\Delta H$ for each of the following reactions:
(a) $2 \mathrm{CH}_{4}(g)+\mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CH}_{3} \mathrm{OH}(g)$
(b) $\mathrm{H}_{2}(g)+\mathrm{Br}_{2}(g) \longrightarrow 2 \mathrm{HBr}(g)$
(c) $2 \mathrm{H}_{2} \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{H}_{2} \mathrm{O}(g)+\mathrm{O}_{2}(g)$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
12:56

Problem 68

Use bond enthalpies (Table $8.4$ ) to estimate the enthalpy change for each of the following reactions:
(a) $\mathrm{C}_{3} \mathrm{H}_{8}(g)+5 \mathrm{O}_{2}(g) \longrightarrow 3 \mathrm{CO}_{2}(g)+4 \mathrm{H}_{2} \mathrm{O}(g)$
(b) $\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}(\mathrm{g})+3 \mathrm{O}_{2}(g) \longrightarrow 2 \mathrm{CO}_{2}(g)+3 \mathrm{H}_{2} \mathrm{O}(g)$
(c) $8 \mathrm{H}_{2} \mathrm{~S}(g) \longrightarrow 8 \mathrm{H}_{2}(g)+\mathrm{S}_{8}(s)$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
04:03

Problem 69

Ammonia is produced directly from nitrogen and hydrogen by using the Haber process. The chemical reaction is
$$\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \longrightarrow 2 \mathrm{NH}_{3}(g)$$
(a) Use bond enthalpies (Table 8.4) to estimate the enthalpy change for the reaction, and tell whether this reaction is exothermic or endothermic. (b) Compare the enthalpy change you calculate in (a) to the true enthalpy change as obtained using $\Delta H_{f}^{\circ}$ values.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
05:40

Problem 70

(a) Use bond enthalpies to estimate the enthalpy change for the reaction of hydrogen with ethene:
$$\mathrm{H}_{2}(g)+\mathrm{C}_{2} \mathrm{H}_{4}(g) \longrightarrow \mathrm{C}_{2} \mathrm{H}_{6}(g)$$
(b) Calculate the standard enthalpy change for this reaction, using heats of formation. Why does this value differ from that calculated in (a)?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
01:30

Problem 71

Given the following bond-dissociation energies, calculate the average bond enthalpy for the $\mathrm{Ti}-\mathrm{Cl}$ bond.
$$
\begin{array}{lc}
& \Delta H(\mathrm{~kJ} / \mathrm{mol}) \\
\hline \mathrm{TiCl}_{4}(g) \longrightarrow \mathrm{TiCl}_{3}(g)+\mathrm{Cl}(g) & 335 \\
\mathrm{TiCl}_{3}(g) \longrightarrow \mathrm{TiCl}_{2}(g)+\mathrm{Cl}(g) & 423 \\
\mathrm{TiCl}_{2}(g) \longrightarrow \mathrm{TiCl}(g)+\mathrm{Cl}(g) & 444 \\
\mathrm{TiCl}(g) \longrightarrow \mathrm{Ti}(g)+\mathrm{Cl}(g) & 519 \\
\hline
\end{array}
$$

Sima Sarker
Sima Sarker
Numerade Educator
06:30

Problem 72

(a) Using average bond enthalpies, predict which of the following reactions will be most exothermic:
(i) $C(g)+2 \mathrm{~F}_{2}(g) \longrightarrow \mathrm{CF}_{4}(g)$
(ii) $\mathrm{CO}(g)+3 \mathrm{~F}_{2} \longrightarrow \mathrm{CF}_{4}(g)+\mathrm{OF}_{2}(g)$
(iii) $\mathrm{CO}_{2}(g)+4 \mathrm{~F}_{2} \longrightarrow \mathrm{CF}_{4}(g)+2 \mathrm{OF}_{2}(g)$
(b) Explain the trend, if any, that exists between reaction exothermicity and the extent to which the carbon atom is bonded to oxygen.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
01:56

Problem 73

How many elements in the periodic table are represented by a Lewis symbol with a single dot? Are all these elements in the same group? Explain.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:59

Problem 74

(a) Explain the following trend in lattice energy: $\mathrm{BeH}_{2}$, $3205 \mathrm{~kJ} / \mathrm{mol} ; \mathrm{MgH}_{2}, 2791 \mathrm{~kJ} / \mathrm{mol} ; \mathrm{CaH}_{2}, 2410 \mathrm{~kJ} / \mathrm{mol} ;$
$\mathrm{SrH}_{2}, 2250 \mathrm{~kJ} / \mathrm{mol} ; \mathrm{BaH}_{2}, 2121 \mathrm{~kJ} / \mathrm{mol}$. (b) The lattice en-
ergy of $\mathrm{ZnH}_{2}$ is $2870 \mathrm{~kJ} / \mathrm{mol}$. Based on the data given in part (a), the radius of the $\mathrm{Zn}^{2+}$ ion is expected to be closest to that of which group $2 \mathrm{~A}$ element?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
08:57

Problem 75

Based on data in Table $8.2$, estimate (within $30 \mathrm{~kJ} / \mathrm{mol}$ ) the lattice energy for (a) LiBr, (b) $\mathrm{CsBr}$, (c) $\mathrm{CaCl}_{2}$.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
01:30

Problem 76

Would you expect AlN to have a lattice energy that is larger or smaller than ScN? Explain.

Nicole Smina
Nicole Smina
Numerade Educator
06:47

Problem 77

From the ionic radii given in Figure $7.8$, calculate the potential energy of a $\mathrm{Ca}^{2+}$ and $\mathrm{O}^{2-}$ ion pair that are just touching (the magnitude of the electronic charge is given on the back inside cover). Calculate the energy of a mole of such pairs. How does this value compare with the lattice energy of $\mathrm{CaO}$ (Table 8.2)? Explain the difference.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
13:09

Problem 78

From Equation $8.4$ and the ionic radii given in Figure 7.8, calculate the potential energy of the following pairs of ions. Assume that the ions are separated by a distance equal to the sum of their ionic radii: (a) $\mathrm{Na}^{+}, \mathrm{Br}^{-}$;
(b) $\mathrm{Rb}^{+}, \mathrm{Br}^{-} ;$ (c) $\mathrm{Sr}^{2+}, \mathrm{S}^{2-}$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:00

Problem 79

(a) How does a polar molecule differ from a nonpolar one? (b) Atoms $X$ and $Y$ have different electronegativities. Will the diatomic molecule $X-Y$ necessarily be polar? Explain. (c) What factors affect the size of the dipole moment of a diatomic molecule?

Nicole Smina
Nicole Smina
Numerade Educator
01:00

Problem 80

Which of the following molecules or ions contain polar bonds: (a) $\mathrm{P}_{4}$, (b) $\mathrm{H}_{2} \mathrm{~S}$, (c) $\mathrm{NO}_{2}^{-}$, (d) $\mathrm{S}_{2}{ }^{2-}$ ?

Nicole Smina
Nicole Smina
Numerade Educator
08:41

Problem 81

To address energy and environmental issues, there is great interest in powering vehicles with hydrogen rather than gasoline. One of the most attractive aspects of the "hydrogen economy" is the fact that in principle the only emission would be water. However, two daunting obstacles must be overcome before this vision can become a reality. First, an economical method of producing hydrogen must be found. Second, a safe, lightweight, and compact way of storing hydrogen must be found. The hydrides of light metals are attractive for hydrogen storage because they can store a high weight percentage of hydrogen in a small volume. One of the most attractive hydrides is $\mathrm{NaAlH}_{4}$, which can release $5.6 \%$ of its mass as $\mathrm{H}_{2}$ upon decomposing to $\mathrm{NaH}(\mathrm{s})$, Al(s), and $\mathrm{H}_{2}(\mathrm{~g})$. NaAlH $_{4}$ possesses both covalent bonds, which hold polyatomic anions together, and ionic bonds. (a) Write a balanced equation for the decomposition of $\mathrm{NaAlH}_{4} .$ (b) Which element in $\mathrm{NaAlH}_{4}$ is the most electronegative? Which one is the least electronegative? (c) Based on electronegativity differences, what do you think is the identity of the polyatomic anion? Draw a Lewis structure for this ion.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
12:11

Problem 82

For the following collection of nonmetallic elements, $\mathrm{O}$, $\mathrm{P}, \mathrm{Te}, \mathrm{I}, \mathrm{B}$, (a) which two would form the most polar single bond? (b) Which two would form the longest single bond? (c) Which two would be likely to form a compound of formula $\mathrm{XY}_{2}$ ? (d) Which combinations of elements would likely yield a compound of empirical formula $\mathrm{X}_{2} \mathrm{Y}_{3} ?$ In each case explain your answer.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
07:45

Problem 83

You and a partner are asked to complete a lab entitled "Oxides of Ruthenium" that is scheduled to extend over two lab periods. The first lab, which is to be completed by your partner, is devoted to carrying out compositional analysis. In the second lab, you are to determine melting points. Upon going to lab you find two unlabeled vials, one containing a soft yellow substance and the other a black powder. You also find the following notes in your partner's notebook-Compound 1: $76.0 \%$ Ru and $24.0 \%$ O (by mass), Compound 2: $61.2 \%$ Ru and $38.8 \% \mathrm{O}$ (by mass).
(a) What is the empirical formula for Compound $1 ?$
(b) What is the empirical formula for Compound 2?
(c) Upon determining the melting points of these two compounds, you find that the yellow compound melts at $25^{\circ} \mathrm{C}$, while the black powder does not melt up to the maximum temperature of your apparatus, $1200^{\circ} \mathrm{C}$. What is the identity of the yellow compound? What is the identity of the black compound? Be sure to use the appropriate naming convention depending upon whether the compound is better described as a molecular or ionic compound.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
09:56

Problem 84

You and a partner are asked to complete a lab entitled "Fluorides of Croup $6 \mathrm{~B}$ Metals" that is scheduled to extend over two lab periods. The first lab, which is to be completed by your partner, is devoted to carrying out compositional analysis. In the second lab, you are to determine melting points. Upon going to lab you find two unlabeled vials, one containing a colorless liquid and the other a green powder. You also find the following notes in your partner's notebook-Compound 1: $47.7 \% \mathrm{Cr}$ and $52.3 \% \mathrm{~F}$ (by mass), Compound 2: $45.7 \% \mathrm{Mo}$ and $54.3 \% \mathrm{~F}$ (by mass). (a) What is the empirical formula for Compound $1 ?$
(b) What is the empirical formula for Compound 2? (c) Upon determining the melting points of these two compounds you find that the colorless liquid solidifies at $18^{\circ} \mathrm{C}$, while the green powder does not melt up to the maximum temperature of your apparatus, $1200^{\circ} \mathrm{C}$. What is the identity of the colorless liquid? What is the identity of the green powder? Be sure to use the appropriate naming convention depending upon whether the compound is better described as a molecular or ionic compound.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
02:18

Problem 85

(a) Triazine, $\mathrm{C}_{3} \mathrm{H}_{3} \mathrm{~N}_{3}$, is like benzene except that in triazine every other $\mathrm{C}-\mathrm{H}$ group is replaced by a nitrogen atom. Draw the Lewis structure(s) for the triazine molecule. (b) Estimate the carbon-nitrogen bond distances in the ring.

Kevin Chimex
Kevin Chimex
Numerade Educator
05:30

Problem 86

Using the electronegativities of $\mathrm{Br}$ and $\mathrm{Cl}$, estimate the partial charges on the atoms in the $\mathrm{Br}-\mathrm{Cl}$ molecule. Using these partial charges and the atomic radii given in Figure $7.7$, estimate the dipole moment of the molecule. The measured dipole moment is $0.57 \mathrm{D}$.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
00:56

Problem 87

Although $\mathrm{I}_{3}^{-}$ is known, $\mathrm{F}_{3}^{-}$ is not. Using Lewis structures, explain why $\mathrm{F}_{3}^{-}$ does not form.

Nicole Smina
Nicole Smina
Numerade Educator
01:21

Problem 88

Calculate the formal charge on the indicated atom in each of the following molecules or ions: (a) the central oxygen atom in $\mathrm{O}_{3},(\mathrm{~b})$ phosphorus in $\mathrm{PF}_{6}^{-}$, (c) nitrogen in $\mathrm{NO}_{2}$, (d) iodine in $\mathrm{ICl}_{3}$, (e) chlorine in $\mathrm{HClO}_{4}$ (hydrogen is bonded to $\mathrm{O}$ ).

James Irizarry
James Irizarry
Numerade Educator
08:20

Problem 89

(a) Determine the formal charge on the chlorine atom in the hypochlorite ion, $\mathrm{ClO}^{-}$, and the perchlorate ion, $\mathrm{ClO}_{4}^{-}$, using resonance structures where the $\mathrm{Cl}$ atom has an octet. (b) What are the oxidation numbers of chlorine in $\mathrm{ClO}^{-}$ and in $\mathrm{ClO}_{4}^{-} ?(\mathrm{c})$ Is it uncommon for the formal charge and the oxidation state to be different? Explain.
(d) Perchlorate is a much stronger oxidizing agent than hypochlorite. Would you expect there to be any relationship between the oxidizing power of the oxyanion and either the oxidation state or the formal charge of chlorine?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
03:44

Problem 90

The following three Lewis structures can be drawn for
$$\begin{aligned}
&\mathrm{N}_{2} \mathrm{O}: \\
&: \mathrm{N} \equiv \mathrm{N}-\ddot{O}: \longleftrightarrow: \ddot{\mathrm{N}}-\mathrm{N} \equiv \mathrm{O}: \longleftrightarrow: \dot{\mathrm{N}}=\mathrm{N}=\ddot{\mathrm{O}}:
\end{aligned}
$$
(a) Using formal charges, which of these three resonance forms is likely to be the most important? (b) The $\mathrm{N}-\mathrm{N}$ bond length in $\mathrm{N}_{2} \mathrm{O}$ is $1.12 \AA$, slightly longer than a typical $\mathrm{N} \equiv \mathrm{N}$ bond; and the $\mathrm{N}-\mathrm{O}$ bond length is $1.19 \AA$, slightly shorter than a typical $\mathrm{N}=\mathrm{O}$ bond. (See Table 8.5.) Rationalize these observations in terms of the resonance structures shown previously and your conclusion for (a).

Nicole Smina
Nicole Smina
Numerade Educator
05:46

Problem 91

An important reaction for the conversion of natural gas to other useful hydrocarbons is the conversion of methane to ethane.
$$
2 \mathrm{CH}_{4}(g) \longrightarrow \mathrm{C}_{2} \mathrm{H}_{6}(g)+\mathrm{H}_{2}(g)
$$
In practice, this reaction is carried out in the presence of oxygen, which converts the hydrogen produced to water.
$$
2 \mathrm{CH}_{4}(g)+\frac{1}{2} \mathrm{O}_{2}(g) \longrightarrow \mathrm{C}_{2} \mathrm{H}_{6}(g)+\mathrm{H}_{2} \mathrm{O}(g)
$$
Use bond enthalpies (Table 8.4) to estimate $\Delta H$ for these two reactions. Why is the conversion of methane to ethane more favorable when oxygen is used?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
12:50

Problem 92

Two compounds are isomers if they have the same chemical formula but a different arrangement of atoms. Use bond enthalpies (Table $8.4)$ to estimate $\Delta H$ for each of the following gas-phase isomerization reactions, and indicate which isomer has the lower enthalpy:

Susan Hallstrom
Susan Hallstrom
Numerade Educator
09:49

Problem 93

With reference to the "Chemistry Put to Work" box on explosives, (a) use bond enthalpies to estimate the enthalpy change for the explosion of $1.00 \mathrm{~g}$ of nitroglycerin. (b) Write a balanced equation for the decomposition of TNT. Assume that, upon explosion, TNT decomposes into $\mathrm{N}_{2}(g), \mathrm{CO}_{2}(g), \mathrm{H}_{2} \mathrm{O}(g)$, and $\mathrm{C}(s)$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
10:04

Problem 94

The "plastic" explosive $\mathrm{C}-4$, often used in action movies, contains the molecule cyclotrimethylenetrinitramine, which is often called RDX (for Royal Demolition eXplosive):
(a) Complete the Lewis structure for the molecule by adding unshared electron pairs where they are needed.
(b) Does the Lewis structure you drew in part (a) have any resonance structures? If so, how many? (c) The molecule causes an explosion by decomposing into $\mathrm{CO}(g)$, $\mathrm{N}_{2}(\mathrm{~g})$, and $\mathrm{H}_{2} \mathrm{O}(g) .$ Write a balanced equation for the decomposition reaction.
(d) With reference to Table $8.4$, which is the weakest type of bond in the molecule?
(e) Use average bond enthalpies to estimate the enthalpy change when $5.0 \mathrm{~g}$ of $\mathrm{RDX}$ decomposes.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
07:12

Problem 95

The bond lengths of carbon-carbon, carbon-nitrogen, carbon-oxygen, and nitrogen-nitrogen single, double, and triple bonds are listed in Table $8.5 .$ Plot bond enthalpy (Table 8.4) versus bond length for these bonds. What do you conclude about the relationship between bond length and bond enthalpy? What do you conclude about the relative strengths of $C-C, C-N, C-O$ and $\mathrm{N}-\mathrm{N}$ bonds?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
09:06

Problem 96

Use the data in Table $8.5$ and the following data: $\mathrm{S}-\mathrm{S}$ distance in $\mathrm{S}_{8}=2.05 \AA ; \mathrm{S}-\mathrm{O}$ distance in $\mathrm{SO}_{2}=1.43 \AA$ to answer the following questions:
(a) Predict the bond length in a S - N single bond.
(b) Predict the bond length in a $\mathrm{S}-\mathrm{O}$ single bond.
(c) Why is the $\mathrm{S}-\mathrm{O}$ bond length in $\mathrm{SO}_{2}$ considerably shorter than your predicted value for the $\mathrm{S}-\mathrm{O}$ single bond?
(d) When elemental sulfur, $\mathrm{S}_{8}$, is carefully oxidized, a compound $\mathrm{S}_{8} \mathrm{O}$ is formed, in which one of the sulfur atoms in the $\mathrm{S}_{8}$ ring is bonded to an oxygen atom. The $\mathrm{S}-\mathrm{O}$ bond length in this compound is $1.48 \AA$. In light of this information, write Lewis structures that can account for the observed $\mathrm{S}-\mathrm{O}$ bond length. Does the sulfur bearing the oxygen in this compound obey the octet rule?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
05:03

Problem 97

The $\mathrm{Ti}^{2+}$ ion is isoelectronic with the Ca atom. (a) Are there any differences in the electron configurations of $\mathrm{Ti}^{2+}$ and Ca? (b) With reference to Figure 6.25, comment on the changes in the ordering of the $4 s$ and $3 d$ subshells in $\mathrm{Ca}$ and $\mathrm{Ti}^{2+}$. (c) Will $\mathrm{Ca}$ and $\mathrm{Ti}^{2+}$ have the same number of unpaired electrons? Explain.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
07:57

Problem 98

(a) Write the chemical equations that are used in calculating the lattice energy of $\mathrm{SrCl}_{2}(s)$ via a Born-Haber cycle.
(b) The second ionization energy of $\operatorname{Sr}(g)$ is $1064 \mathrm{~kJ} / \mathrm{mol}$. Use this fact along with data in Appendix C, Figure 7.12, Figure $7.14$, and Table $8.2$ to calculate $\Delta H_{f}^{\circ}$ for $\operatorname{SrCl}_{2}(s)$.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
05:09

Problem 99

The electron affinity of oxygen is $-141 \mathrm{~kJ} / \mathrm{mol}$, corresponding to the reaction
$$
\mathrm{O}(g)+\mathrm{e}^{-}-\rightarrow \mathrm{O}^{-}(g)
$$
The lattice energy of $\mathrm{K}_{2} \mathrm{O}(s)$ is $2238 \mathrm{~kJ} / \mathrm{mol}$. Use these data along with data in Appendix $C$ and Figure $7.12$ to calculate the "second electron affinity" of oxygen, corresponding to the reaction
$$
\mathrm{O}^{-}(g)+\mathrm{e}^{-} \longrightarrow \mathrm{O}^{2-}(g)
$$

James Irizarry
James Irizarry
Numerade Educator
13:37

Problem 100

The reaction of indium, In, with sulfur leads to three binary compounds, which we will assume to be purely ionic. The three compounds have the following properties:
$$
\begin{array}{lll}
\hline \text { Compound } & \text { Mass \% In } & \text { Melting Point }\left({ }^{\circ} \mathrm{C}\right) \\
\hline \text { A } & 87.7 & 653 \\
\text { B } & 78.2 & 692 \\
\text { C } & 70.5 & 1050 \\
\hline
\end{array}
$$
(a) Determine the empirical formulas of compounds A, B, and C. (b) Give the oxidation state of In in each of the three compounds. (c) Write the electron configuration for the In ion in each compound. Do any of these configurations correspond to a noble-gas configuration? (d) In which compound is the ionic radius of In expected to be smallest? Explain. (e) The melting point of ionic compounds often correlates with the lattice energy. Explain the trends in the melting points of compounds $A, B$, and C in these terms.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
24:13

Problem 101

One scale for electronegativity is based on the concept that the electronegativity of any atom is proportional to the ionization energy of the atom minus its electron affinity: electronegativity $=k(\mathrm{IE}-\mathrm{EA})$, where $k$ is a proportionality constant. (a) How does this definition explain why the electronegativity of $\mathrm{F}$ is greater than that of $C l$ even though $C l$ has the greater electron affinity? (b) Why are both ionization energy and electron affinity relevant to the notion of electronegativity? (c) By using data in Chapter 7 , determine the value of $k$ that would lead to an electronegativity of $4.0$ for $\mathrm{F}$ under this definition. (d) Use your result from part (c) to determine the electronegativities of $\mathrm{Cl}$ and $\mathrm{O}$ using this scale. Do these values follow the trend shown in Figure 8.6?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
04:54

Problem 102

The compound chloral hydrate, known in detective stories as knockout drops, is composed of $14.52 \% \mathrm{C}$, $1.83 \% \mathrm{H}, 64.30 \% \mathrm{Cl}$, and $19.35 \%$ O by mass and has a molar mass of $165.4 \mathrm{~g} / \mathrm{mol}$. (a) What is the empirical formula of this substance? (b) What is the molecular formula of this substance? (c) Draw the Lewis structure of the molecule, assuming that the $\mathrm{Cl}$ atoms bond to a single $C$ atom and that there are a $C-C$ bond and two $\mathrm{C}-\mathrm{O}$ bonds in the compound.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
07:20

Problem 103

Barium azide is $62.04 \%$ Ba and $37.96 \%$ N. Each azide ion has a net charge of $1-$. (a) Determine the chemical formula of the azide ion. (b) Write three resonance structures for the azide ion. (c) Which structure is most important? (d) Predict the bond lengths in the ion.

James Irizarry
James Irizarry
Numerade Educator
04:54

Problem 104

Acetylene $\left(\mathrm{C}_{2} \mathrm{H}_{2}\right)$ and nitrogen $\left(\mathrm{N}_{2}\right)$ both contain a triple bond, but they differ greatly in their chemical properties. (a) Write the Lewis structures for the two substances. (b) By referring to Appendix $C$, look up the enthalpies of formation of acetylene and nitrogen and compare their reactivities. (c) Write balanced chemical equations for the complete oxidation of $\mathrm{N}_{2}$ to form $\mathrm{N}_{2} \mathrm{O}_{5}(g)$ and of acetylene to form $\mathrm{CO}_{2}(g)$ and $\mathrm{H}_{2} \mathrm{O}(g)$
(d) Calculate the enthalpy of oxidation per mole of $\mathrm{N}_{2}$ and $\mathrm{C}_{2} \mathrm{H}_{2}$ (the enthalpy of formation of $\mathrm{N}_{2} \mathrm{O}_{5}(g)$ is $11.30 \mathrm{~kJ} / \mathrm{mol}$ ). How do these comparative values relate to your response to part (b)? Both $\mathrm{N}_{2}$ and $\mathrm{C}_{2} \mathrm{H}_{2}$ possess triple bonds with quite high bond enthalpies (Table 8.4). What aspect of chemical bonding in these molecules or in the oxidation products seems to account for the difference in chemical reactivities?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
14:37

Problem 105

Under special conditions, sulfur reacts with anhydrous liquid ammonia to form a binary compound of sulfur and nitrogen. The compound is found to consist of $69.6 \% \mathrm{~S}$ and $30.4 \%$ N. Measurements of its molecular mass yield a value of $184.3 \mathrm{~g} \mathrm{~mol}^{-1}$. The compound occasionally detonates on being struck or when heated rapidly. The sulfur and nitrogen atoms of the molecule are joined in a ring. All the bonds in the ring are of the same length. (a) Calculate the empirical and molecular formulas for the substance. (b) Write Lewis structures for the molecule, based on the information you are given. (Hint: You should find a relatively small number of dominant Lewis structures.) (c) Predict the bond distances between the atoms in the ring. (Note: The $\mathrm{S}-\mathrm{S}$ distance in the $\mathrm{S}_{8}$ ring is $2.05 \AA$.) (d) The enthalpy of formation of the compound is estimated to be $480 \mathrm{~kJ} \mathrm{~mol}^{-1}$. $\Delta H_{f}^{\circ}$ of $S(g)$ is $222.8 \mathrm{~kJ} \mathrm{~mol}^{-1}$. Estimate the average bond enthalpy in the compound.

Susan Hallstrom
Susan Hallstrom
Numerade Educator
04:25

Problem 106

A common form of elemental phosphorus is the tetrahedral $\mathrm{P}_{4}$ molecule, where all four phosphorus atoms are equivalent: At room temperature phosphorus is a solid. (a) Do you think there are any unshared pairs of electrons in the $\mathrm{P}_{4}$ molecule? (b) How many $\mathrm{P}-\mathrm{P}$ bonds are there in the molecule? (c) Can you draw a Lewis structure for a linear $\mathrm{P}_{4}$ molecule that satisfies the octet rule? (d) Using formal charges, what can you say about the stability of the linear molecule vs. that of the tetrahedral molecule?

Susan Hallstrom
Susan Hallstrom
Numerade Educator
12:04

Problem 107

Consider benzene $\left(\mathrm{C}_{6} \mathrm{H}_{6}\right)$ in the gas phase. (a) Write the reaction for breaking all the bonds in $\mathrm{C}_{6} \mathrm{H}_{6}(g)$, and use data in Appendix $C$ to determine the enthalpy change for this reaction. (b) Write a reaction that corresponds to breaking all the carbon-carbon bonds in $\mathrm{C}_{6} \mathrm{H}_{6}(g) .$ (c) By combining your answers to parts (a) and (b) and using the average bond enthalpy for $\mathrm{C}-\mathrm{H}$ from Table $8.4$, calculate the average bond enthalpy for the carbon-carbon bonds in $\mathrm{C}_{6} \mathrm{H}_{6}(g) .$ (d) Comment on your answer from part (c) as compared to the values for $C-C$ single bonds and $\mathrm{C}=\mathrm{C}$ double bonds in Table $8.4 .$

Susan Hallstrom
Susan Hallstrom
Numerade Educator
12:48

Problem 108

Average bond enthalpies are generally defined for gasphase molecules. Many substances are liquids in their standard state. $\mathrm{OB}$ (Section 5.7) By using appropriate thermochemical data from Appendix C, calculate average bond enthalpies in the liquid state for the following bonds, and compare these values to the gas-phase values given in Table 8.4: (a) $\mathrm{Br}-\mathrm{Br}$, from $\mathrm{Br}_{2}(l) ;$ (b) $\mathrm{C}-\mathrm{Cl}$, from $\mathrm{CCl}_{4}(l) ;$ (c) $\mathrm{O}-\mathrm{O}$, from $\mathrm{H}_{2} \mathrm{O}_{2}(l)$ (assume that the $\mathrm{O}-\mathrm{H}$ bond enthalpy is the same as in the gas phase). (d) What can you conclude about the process of breaking bonds in the liquid as compared to the gas phase? Explain the difference in the $\Delta H$ values between the two phases.

Susan Hallstrom
Susan Hallstrom
Numerade Educator