Answer: Draw the Lewis structures for each of the

For each of these Lewis symbols, indicate the group in the periodic table in which the element X belongs: [Section 8.1] (a) X (b) X (c) X

Illustrated are four ions A, B, X, and Y showing their relative ionic radii. The ions shown in red carry positive charges: The potential energy of two interacting charges m = Qr [8.10] The dipole moment of two charges of equal magnitude but opposite sign, separated by a distance r Formal charge = valence electrons - [8.11] 1 2 1bonding electrons2 - nonbonding electrons The definition of formal charge Hrxn = g1bond enthalpies of bonds broken2 - [8.12] g1bond enthalpies of bonds formed2 The enthalpy change as a function of bond enthalpies for reactions involving gas-phase molecules a 2+ charge for A and a 1+ charge for 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 the combinations in part (a), which leads to the ionic compound having the largest lattice energy? [Section 8.2]

A portion of a two-dimensional slab of NaCl(s) is shown here (see Figure 8.3) in which the ions are numbered. (a) Which colored balls must represent sodium ions? (b) Which colored balls must represent chloride ions? (c) Consider ion 5. How many attractive electrostatic interactions are shown for it? (d) Consider ion 5. How many repulsive interactions are shown for it? (e) Is the sum of the attractive interactions in part (c) larger or smaller than the sum of the repulsive interactions in part (d)? (f) If this pattern of ions was extended indefinitely in two dimensions, would the lattice energy be positive or negative? [Section 8.2] 1 2 3 7 8 9

The orbital diagram that follows 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] 4d

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

Incomplete Lewis structures for the nitrous acid molecule, HNO2, and the nitrite ion, NO2 -, are shown here. (a) Complete each Lewis structure by adding electron pairs as needed. (b) Is the formal charge on N the same or different in these two species? (c) Would either HNO2 or NO2 - be expected to exhibit resonance? (d) Would you expect the NO bond in HNO2 to be longer, shorter, or the same length as the NO bonds in NO2 -? [Sections 8.5 and 8.6] HONO ONO

The partial Lewis structure that follows 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 carboncarbon bonds are labeled 1, 2, and 3. (a) How many hydrogen atoms are in the molecule? (b) Rank the carboncarbon bonds in order of increasing bond length. (c) Rank the carboncarbon bonds in order of increasing bond enthalpy. [Sections 8.3 and 8.8] C C C C 1 2 3

Consider the Lewis structure for the polyatomic oxyanion shown here, where X is an element from the third period 1Na - Ar2. 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, X; (b) determine the formal charge of the central atom, X; (c) draw a Lewis structure that makes the formal charge on the central atom equal to zero. [Sections 8.5, 8.6, and 8.7] O O O O X

(a) True or false: An elements number of valence electrons is the same as its atomic number.

(b) How many valence electrons does a nitrogen atom possess?

(c) An atom has the electron configuration 1s 2 2s 2 2p6 3s 2 3p2 . How many valence electrons does the atom have

(a) True or false: The hydrogen atom is most stable when it has a full octet of electrons

(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 1s 2 2s 2 2p3 , how many electrons must it gain to achieve an octet?

Consider the element silicon, Si. (a) Write its electron configuration. (b) How many valence electrons does a silicon atom have? (c) Which subshells hold the valence electrons?

(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 4B. Write the electron configuration for Hf.

(c) 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?

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

What is the Lewis symbol for each of the following atoms or ions? (a) K, (b) As, (c) Sn2 + , (d) N3 - .

(a) Using Lewis symbols, diagram the reaction between magnesium and oxygen atoms to give the ionic substance MgO.

(b) How many electrons are transferred?

(c) Which atom loses electrons in the reaction?

(a) Use Lewis symbols to represent the reaction that occurs between Ca and F atoms

(b) What is the chemical formula of the most likely product?

(c) How many electrons are transferred?

(d) Which atom loses electrons in the reaction?

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

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.

Write the electron configuration for each of the following ions, and determine which ones possess noble-gas configurations: (a) Sr 2+, (b) Ti 2+, (c) Se 2-, (d) Ni 2+, (e) Br -, (f) Mn3 + .

Write electron configurations for the following ions, and determine which have noble-gas configurations: (a) Cd2+, (b) P3-, (c) Zr 4+, (d) Ru3+, (e) As 3-, (f) Ag+.

(a) Is lattice energy usually endothermic or exothermic?

(b) Write the chemical equation that represents the process of lattice energy for the case of NaCl.

(c) Would you expect salts like NaCl, which have singly-charged ions, to have larger or smaller lattice energies compared to salts like CaO which are composed of doubly-charged ions?

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 NaCl and KF are given in Table 8.2. Based on the lattice energies, would you expect the NaCl or the KF distance to be longer? (b) Use the ionic radii given in Figure 7.8 to estimate the NaCl and KF distances.

The ionic substances NaF, CaO, and ScN are isoelectronic (they have the same number of electrons). Examine the lattice energies for these substances in Table 8.2. Make a graph of lattice energy on the vertical axis versus the charge on the cation on the horizontal axis. (a) What is the slope of the line? (b) Make a graph of lattice energy on the vertical axis versus the square of the cation charge on the horizontal axis. What is the slope of this line? (c) Compare how well the data points fall on a line for the graphs in (a) and (b). Which trend is more linear, lattice energy versus cation charge or lattice energy versus cation charge squared? (d) Predict the lattice energy for the compound TiC, if we consider the carbon to have a 4- charge.

(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) Arrange the following substances not listed in Table 8.2 according to their expected lattice energies, listing them from lowest lattice energy to the highest: MgS, KI, GaN, LiBr.

Consider the ionic compounds KF, NaCl, NaBr, and LiCl. (a) Use ionic radii (Figure 7.8) to estimate the cationanion distance for each compound. (b) Based on your answer to part (a), arrange these four compounds in order of decreasing lattice energy. (c) Check your predictions in part (b) with the experimental values of lattice energy from Table 8.2. Are the predictions from ionic radii correct?

Which of the following trends in lattice energy is due to differences in ionic radii? (a) NaCl 7 RbBr 7 CsBr, (b) BaO 7 KF, (c) SrO 7 SrCl2.

Energy is required to remove two electrons from Ca to form \(\mathrm{Ca}^{2+}\), and energy is required to add two electrons to \(\mathrm{O}\) to form \(\mathrm{O}^{2-}\). Yet \(\mathrm{CaO}\) is stable relative to the free elements. Which statement is the best explanation? (a) The lattice energy of \(\mathrm{CaO}\) is large enough to overcome these processes. (b) \(\mathrm{CaO}\) is a covalent compound, and these processes are irrelevant. (c) \(\mathrm{CaO}\) has a higher molar mass than either \(\mathrm{Ca}\) or \(\mathrm{O}\). (d) The enthalpy of formation of \(\mathrm{CaO}\) is small. (e) \(\mathrm{CaO}\) is stable to atmospheric conditions.

List the individual steps used in constructing a BornHaber cycle for the formation of BaI2 from the elements. Which of the steps would you expect to be exothermic?

Use data from Appendix C, Figure 7.10, and Figure 7.12 to calculate the lattice energy of RbCl.

(a) Based on the lattice energies of MgCl2 and SrCl2 given in Table 8.2, what is the range of values that you would expect for the lattice energy of CaCl2?

(b) Using data from Appendix C, Figure 7.11, Figure 7.13 and the value of the second ionization energy for Ca, 1145 kJ>mol, calculate the lattice energy of CaCl2.

(a) State whether the bonding in each compound is likely to be covalent or not: (i) iron, (ii) sodium chloride, (iii) water, (iv) oxygen, (v) argon.

(b) A substance XY, formed from two different elements, boils at -33 C. Is XY likely to be a covalent or an ionic substance?

Which of these elements are unlikely to form covalent bonds? S, H, K, Ar, Si.

Using Lewis symbols and Lewis structures, diagram the formation of SiCl4 from Si and Cl atoms, showing valence-shell electrons. (a) How many valence electrons does Si have initially? (b) How many valence electrons does each Cl have initially? (c) How many valence electrons surround the Si in the SiCl4 molecule? (d) How many valence electrons surround each Cl in the SiCl4 molecule? (e) How many bonding pairs of electrons are in the SiCl4 molecule?

Use Lewis symbols and Lewis structures to diagram the formation of PF3 from P and F atoms, showing valence-shell electrons. (a) How many valence electrons does P have initially? (b) How many valence electrons does each F have initially? (c) How many valence electrons surround the P in the PF3 molecule? (d) How many valence electrons surround each F in the PF3 molecule? (e) How many bonding pairs of electrons are in the PF3 molecule?

(a) Construct a Lewis structure for O2 in which each atom achieves an octet of electrons.

(b) How many bonding electrons are in the structure?

(c) Would you expect the OO bond in O2 to be shorter or longer than the OO bond in compounds that contain an OO single bond? Explain.

(a) Construct a Lewis structure for hydrogen peroxide, H2O2, in which each atom achieves an octet of electrons.

(b) How many bonding electrons are between the two oxygen atoms?

(c) Do you expect the OO bond in H2O2 to be longer or shorter than the OO bond in O2? Explain.

Which of the following statements about electronegativity is false? (a) Electronegativity is the ability of an atom in a molecule to attract electron density toward itself. (b) Electronegativity is the same thing as electron affinity. (c) The numerical values for electronegativity have no units. (d) Fluorine is the most electronegative element. (e) Cesium is the least electronegative element.

(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) True or false: The most easily ionizable elements are the most electronegative.

Using only the periodic table as your guide, select the most electronegative atom in each of the following sets: (a) Na, Mg, K, Ca; (b) P, S, As, Se; (c) Be, B, C, Si; (d) Zn, Ge, Ga, As.

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

Which of the following bonds are polar? (a) BF, (b) ClCl, (c) SeO, (d) HI. Which is the more electronegative atom in each polar bond?

Arrange the bonds in each of the following sets in order of increasing polarity: (a) CF, OF, BeF; (b) OCl, SBr, CP; (c) CS, BF, NO.

(a) From the data in Table 8.3, calculate the effective charges on the H and Br atoms of the HBr molecule in units of the electronic charge, e.

(b) If you were to put HBr under very high pressure, so its bond length decreased significantly, would its dipole moment increase, decrease, or stay the same, if you assume that the effective charges on the atoms do not change?

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

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) SiF4 and LaF3, (b) FeCl2 and ReCl6, (c) PbCl4 and RbCl

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) TiCl4 and CaF2, (b) ClF3 and VF3, (c) SbCl5 and AlF3.

Draw Lewis structures for the following: (a) SiH4, (b) CO, (c) SF2, (d) H2SO4 (H is bonded to O), (e) ClO2 -, (f) NH2OH.

Write Lewis structures for the following: (a) H2CO (both H atoms are bonded to C), (b) H2O2, (c) C2F6 (contains a CC bond), (d) AsO3 3 - , (e) H2SO3 (H is bonded to O), (f) NH2Cl.

Which one of these statements about formal charge is true? (a) Formal charge is the same as oxidation number. (b) To draw the best Lewis structure, you should minimize formal charge. (c) Formal charge takes into account the different electronegativities of the atoms in a molecule. (d) Formal charge is most useful for ionic compounds. (e) Formal charge is used in calculating the dipole moment of a diatomic molecule.

(a) Draw the dominant Lewis structure for the phosphorus trifluoride molecule, PF3.

(b) Determine the oxidation numbers of the P and F atoms

(c) Determine the formal charges of the P and F atoms.

Write Lewis structures that obey the octet rule for each of the following, and assign oxidation numbers and formal charges to each atom: (a) OCS, (b) SOCl2 (S is the central atom), (c) BrO3 -, (d) HClO2 (H is bonded to O).

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) SO2, (b) SO3, (c) SO3 2 - . (d) Arrange these molecules/ions in order of increasing SO bond length

(a) Draw the best Lewis structure(s) for the nitrite ion, NO2 -

(b) With what allotrope of oxygen is it isoelectronic?

(c) What would you predict for the lengths of the bonds in NO2 - relative to NO single bonds and double bonds?

Consider the formate ion, HCO2 -, which is the anion formed when formic acid loses an H+ ion. The H and the two O atoms are bonded to the central C atom. (a) Draw the best Lewis structure(s) for this ion. (b) Are resonance structures needed to describe the structure? (c) Would you predict that the CO bond lengths in the formate ion would be longer or shorter relative to those in CO2?

Predict the ordering, from shortest to longest, of the bond lengths in CO, CO2, and CO3 2 - .

Based on Lewis structures, predict the ordering, from shortest to longest, of NO bond lengths in NO+, NO2 -, and NO3 -

(a) Do the CC bond lengths in benzene alternate shortlong-short-long around the ring? Why or why not?

(b) Are CC bond lengths in benzene shorter than CC single bonds?

(c) Are CC bond lengths in benzene shorter than CC double bonds?

Mothballs are composed of naphthalene, C10H8, a molecule that consists of two six-membered rings of carbon fused along an edge, as shown in this incomplete Lewis structure: C H H H C C H H C C C H H H C C C C (a) Draw all of the resonance structures of naphthalene. How many are there? (b) Do you expect the CC bond lengths in the molecule to be similar to those of CC single bonds, CC double bonds, or intermediate between CC single and CC double bonds? (c) Not all of the CC bond lengths in naphthalene are equivalent. Based on your resonance structures, how many CC bonds in the molecule do you expect to be shorter than the others?

Indicate whether each statement is true or false: (a) The octet rule is based on the fact that filling in all s and p valence electrons in a shell gives eight electrons. (b) The Si in SiH4 does not follow the octet rule because hydrogen is in an unusual oxidation state. (c) Boron compounds are frequent exceptions to the octet rule because they have too few electrons surrounding the boron. (d) Compounds in which nitrogen is the central atom are frequent exceptions to the octet rule because they have too many electrons surrounding the nitrogen.

Fill in the blank with the appropriate numbers for both electrons and bonds (considering that single bonds are counted as one, double bonds as two, and triple bonds as three). (a) Fluorine has ____ valence electrons and makes ____ bond(s) in compounds. (b) Oxygen has ____ valence electrons and makes ____ bond(s) in compounds. (c) Nitrogen has ____ valence electrons and makes ____ bond(s) in compounds. (d) Carbon has ____ valence electrons and makes____ bond(s) in compounds.

Draw the dominant Lewis structures for these chlorineoxygen molecules/ions: ClO, ClO-, ClO2 -, ClO3 -, ClO4 -. Which of these do not obey the octet rule?

For elements in the third row of the periodic table and beyond, the octet rule is often not obeyed. A friend of yours says this is because these heavier elements are more likely to make double or triple bonds. Another friend of yours says that this is because the heavier elements are larger and can make bonds to more than four atoms at a time. Which friend is most correct?

Draw the Lewis structures for each of the following ions or molecules. Identify those in which the octet rule is not obeyed; state which atom in each compound does not follow the octet rule; and state, for those atoms, how many electrons surround these atoms: (a) PH3, (b) AlH3, (c) N3 -, (d) CH2Cl2, (e) SnF6.

Draw the Lewis structures for each of the following molecules or ions. Identify instances where the octet rule is not obeyed; state which atom in each compound does not follow the octet rule; and state how many electrons surround these atoms: (a) NO, (b) BF3, (c) ICl2 -, (d) OPBr3 (the P is the central atom), (e) XeF4

In the vapor phase, BeCl2 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 structures are possible that satisfy the octet rule? (c) On the basis of the formal charges, which Lewis structure is expected to be dominant for BeCl2?

(a) Describe the molecule xenon trioxide, XeO3, using four possible Lewis structures, one each with zero, one, two, or three XeO double bonds.

(b) Do any of these resonance structures satisfy the octet rule for every atom in the molecule?

(c) Do any of the four Lewis structures have multiple resonance structures? If so, how many resonance structures do you find?

(d) Which of the Lewis structures in (a) yields the most favorable formal charges for the molecule?

Consider the following statement: For some molecules and ions, a Lewis structure that satisfies the octet rule does not lead to the lowest formal charges, and a Lewis structure that leads to the lowest formal charges does not satisfy the octet rule. Illustrate this statement using the hydrogen sulfite ion, HSO3 -, as an example (the H atom is bonded to one of the O atoms).

Using Table 8.4, estimate H for each of the following gasphase reactions (note that lone pairs on atoms are not shown): H H H H H O O H O C O H H H H H C 2 Cl N Cl + 3 Cl Cl Cl C C H H H H H H C H C H H H H C C C N N N C N (a) (b) (c)

Using Table 8.4, estimate H for the following gas-phase reactions: Br C Br Br H + Cl Cl Cl Cl + H Cl H N N H H + H H S C H H C S H + 2 H H H Br C H H C Br + H H 2 H S H 2 H N H Cl Br C Br Br Br (a) Cl (b) (c)

Using Table 8.4, estimate H for each of the following reactions: (a) 2 CH41g2 + O21g2 2 CH3OH1g2 (b) H21g2 + Br21l2 2 HBr1g2 (c) 2 H2O21g2 2 H2O1g2 + O21g2

Use Table 8.4 to estimate the enthalpy change for each of the following reactions: (a) H2CO1g2 + HCl1g2 H3COCl1g2 (b) H2O21g2+2 CO1g2 H21g2+2 CO21g2 (c) 3 H2CCH21g2 C6H121g2 (the six carbon atoms form a six-membered ring with two H atoms on each C atom)

Ammonia is produced directly from nitrogen and hydrogen by using the Haber process, which is perhaps the most widely used industrial chemical reaction on Earth. The chemical reaction is N21g2 + 3 H21g2 2 NH31g2 (a) Use Table 8.4 to estimate the enthalpy change for the reaction. Is it exothermic or endothermic? (b) Calculate the enthalpy change as obtained using Hf values.

(a) Use bond enthalpies to estimate the enthalpy change for the reaction of hydrogen with ethylene: H21g2 + C2H41g2 C2H61g2

(b) Calculate the standard enthalpy change for this reaction, using heat of formation.

Given the following bond-dissociation energies, calculate the average bond enthalpy for the TiCl bond. H1kJ>mol2 TiCl41g2 TiCl31g2 + Cl1g2 335 TiCl31g2 TiCl21g2 + Cl1g2 423 TiCl21g2 TiCl1g2 + Cl1g2 444 TiCl1g2 Ti1g2 + Cl1g2 519

(a) Using average bond enthalpies, predict which of the following reactions will be most exothermic: (i) C1g2 + 2 F21g2 CF41g2 (ii) CO1g2 + 3 F21g2 CF41g2 + OF21g2 (iii) CO21g2 + 4 F21g2 CF41g2 + 2 OF21g2

(b) Make a graph that plots the reaction enthalpy you calculated on the vertical axis versus oxidation state of carbon on the horizontal axis. Draw the best-fit line through your data points. Is the slope of this line negative or positive? Does this mean that as the oxidation state of the carbon increases, the reaction with elemental fluorine becomes more or less exothermic?

(c) Predict the reaction enthalpy for the reaction of carbonate ion with fluorine from your graph, and compare to what you predict using average bond enthalpies.

How many elements in the periodic table are represented by a Lewis symbol with a single dot? Which groups are they in?

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) Na +, Br -; (b) Rb+, Br -; (c) Sr 2 + , S2 - .

(a) Consider the lattice energies for the following compounds: BeH2, 3205 kJ>mol; MgH2, 2791 kJ>mol; CaH2, 2410 kJ>mol; SrH2, 2250 kJ>mol; BaH2, 2121 kJ>mol. Plot lattice energy versus cation radius for these compounds. If you draw a line through your points, is the slope negative or positive? Explain.

(b) The lattice energy of ZnH2 is 2870 kJ>mol. Based on the data given in part (a), the radius of the Zn2 + ion is expected to be closest to that of which group 2A element?

Based on data in Table 8.2, estimate (within 30 kJ>mol) the lattice energy for (a) LiBr, (b) CsBr, (c) CaCl2

An ionic substance of formula MX has a lattice energy of 6 * 103 kJ>mol. Is the charge on the ion M likely to be 1+, 2+, or 3+? Explain.

From the ionic radii given in Figure 7.8, calculate the potential energy of a Ca 2+ and O2- ion pair that is 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 CaO (Table 8.2)? Explain

Construct a BornHaber cycle for the formation of the hypothetical compound NaCl2, where the sodium ion has a 2+ charge (the second ionization energy for sodium is given in Table 7.2). (a) How large would the lattice energy need to be for the formation of NaCl2 to be exothermic? (b) If we were to estimate the lattice energy of NaCl2 to be roughly equal to that of MgCl2 (2326 kJ>mol from Table 8.2), what value would you obtain for the standard enthalpy of formation, Hf , of NaCl2?

A classmate of yours is convinced that he knows everything about electronegativity. (a) In the case of atoms X and Y having different electronegativities, he says, the diatomic molecule XY must be polar. Is your classmate correct? (b) Your classmate says that the farther the two atoms are apart in a bond, the larger the dipole moment will be. Is your classmate correct?

Consider the collection of nonmetallic elements O, P, Te, I and 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 XY2? (d) Which combinations of elements would likely yield a compound of empirical formula X2Y3?

The substance chlorine monoxide, ClO(g), is important in atmospheric processes that lead to depletion of the ozone layer. The ClO molecule has an experimental dipole moment of 1.24 D, and the ClO bond length is 1.60 . (a) Determine the magnitude of the charges on the Cl and O atoms in units of the electronic charge, e. (b) Based on the electronegativities of the elements, which atom would you expect to have a partial negative charge in the ClO molecule? (c) Using formal charges as a guide, propose the dominant Lewis structure for the molecule. (d) The anion ClO- exists. What is the formal charge on the Cl for the best Lewis structure for ClO-?

(a) Using the electronegativities of Br and Cl, estimate the partial charges on the atoms in the BrCl molecule.

(b) Using these partial charges and the atomic radii given in Figure 7.8, estimate the dipole moment of the molecule

(c) The measured dipole moment of BrCl is 0.57 D. If you assume the bond length in BrCl is the sum of the atomic radii, what are the partial charges on the atoms in BrCl using the experimental dipole moment?

A major challenge in implementing the "hydrogen economy" is finding a safe, lightweight, and compact way of storing hydrogen for use as a fuel. The hydrides of light metals are attractive for hydrogen storage because they can store a high weight percentage of hydrogen in a small volume. For example, \(\mathrm{NaAlH}_{4}\) can release \(5.6 \%\) of its mass as \(\mathrm{H}_{2}\) upon decomposing to \(\mathrm{NaH}(s), \mathrm{Al}(s)\), and \(\mathrm{H}_{2}(g)\). \(\mathrm{NaAlH}_{4}\) 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, predict the identity of the polyatomic anion. Draw a Lewis structure for this ion. (d) What is the formal charge on hydrogen in the polyatomic ion?

Although I3 - is known, F3 - is not. Which statement is the most correct explanation? (a) Iodine is more likely to be electron-deficient. (b) Fluorine is too small to accommodate three nonbonding electron pairs and two bonding electron pairs. (c) Fluorine is too electronegative to form anions. (d) I2 is known but F2 is not. (e) Iodine has a larger electron affinity than fluorine

Calculate the formal charge on the indicated atom in each of the following molecules or ions: (a) the central oxygen atom in O3, (b) phosphorus in PF6 -, (c) nitrogen in NO2, (d) iodine in ICl3, (e) chlorine in HClO4 (hydrogen is bonded to O).

(a) Determine the formal charge on the chlorine atom in the hypochlorite ion, ClO-, and the perchlorate ion, ClO4 -, using resonance structures where the Cl atom has an octet.

(b) What are the oxidation numbers of chlorine in ClO- and in ClO4 -?

(c) Perchlorate is a much stronger oxidizing agent than hypochlorite. Suggest an explanation.

The following three Lewis structures can be drawn for N2O: N N O N N O N N O (a) Using formal charges, which of these three resonance forms is likely to be the most important? (b) The NN bond length in N2O is 1.12 , slightly longer than a typical NN bond; and the NO bond length is 1.19 , slightly shorter than a typical bond (see Table 8.5). Based on these data, which resonance structure best represents N2O?

(a) Triazine, C3H3N3, is like benzene except that in triazine every other CH group is replaced by a nitrogen atom. Draw the Lewis structure(s) for the triazine molecule.

(b) Estimate the carbonnitrogen bond distances in the ring.

Ortho-dichlorobenzene, C6H4Cl2, is obtained when two of the adjacent hydrogen atoms in benzene are replaced with Cl atoms. A skeleton of the molecule is shown here. (a) Complete a Lewis structure for the molecule using bonds and electron pairs as needed. (b) Are there any resonance structures for the molecule? If so, sketch them. (c) Are the resonance structures in (a) and (b) equivalent to one another as they are in benzene?

Consider the hypothetical molecule BAB. Are the following statements true or false? (a) This molecule cannot exist. (b) If resonance was important, the molecule would have identical AB bond lengths.

An important reaction for the conversion of natural gas to other useful hydrocarbons is the conversion of methane to ethane. 2 CH41g2 C2H61g2 + H21g2 In practice, this reaction is carried out in the presence of oxygen, which converts the hydrogen produced into water. 2 CH41g2 + 1 2O21g2 C2H61g2 + H2O1g2 Use Table 8.4 to estimate H for these two reactions. Why is the conversion of methane to ethane more favorable when oxygen is used?

Two compounds are isomers if they have the same chemical formula but different arrangements of atoms. Use Table 8.4 to estimate H for each of the following gas-phase isomerization reactions and indicate which isomer has the lower enthalpy. Ethanol Dimethyl ether H C H H H C H H C H H O H O C H H H (a) H C H O O Ethylene oxide H H C H H C H C H Acetaldehyde (b) C C C C H H H H H H H H H H H H C H C H C H C C C H Cyclopentene Pentadiene (c) H H H C N C H H H C C N Methyl isocyanide Acetonitrile (d

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 g of nitroglycerin. (b) Write a balanced equation for the decomposition of TNT. Assume that, upon explosion, TNT decomposes into N21g2, CO21g2, H2O1g2, and C(s).

The plastic explosive C-4, often used in action movies, contains the molecule cyclotrimethylenetrinitramine, which is often called RDX (for Royal Demolition eXplosive): O N N C C C N H H N H H H H N O O O Cyclotrimethylenetrinitramine (RDX) N O O (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 CO1g2, N21g2, and H2O1g2. 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 g of RDX decomposes.

The bond lengths of carboncarbon, carbonnitrogen, carbonoxygen, and nitrogennitrogen single, double, and triple bonds are listed in Table 8.5. Plot bond enthalpy (Table 8.4) versus bond length for these bonds (as in Figure 8.17). (a) Is this statement true: The longer the bond, the stronger the bond? (b) Order the relative strengths of CC, CN, CO, and NN bonds from weakest to strongest. (c) From your graph for the carboncarbon bond in part (a), estimate the bond enthalpy of the hypothetical CC quadruple bond

The Ti 2 + ion is isoelectronic with the Ca atom. (a) Write the electron configurations of Ti 2 + and Ca. (b) Calculate the number of unpaired electrons for Ca and for Ti 2 + . (c) What charge would Ti have to be isoelectronic with Ca 2 + ?

(a) Write the chemical equations that are used in calculating the lattice energy of SrCl21s2 via a BornHaber cycle.

(b) The second ionization energy of Sr(g) is 1064 kJ>mol. Use this fact along with data in Appendix C, Figure 7.10, Figure 7.12, and Table 8.2 to calculate Hf for SrCl21s2.

The electron affinity of oxygen is -141 kJ>mol, corresponding to the reaction O21g2 + e- O-1g2 The lattice energy of K2O1s2 is 2238 kJ>mol. Use these data along with data in Appendix C and Figure 7.10 to calculate the second electron affinity of oxygen, corresponding to the reaction O-1g2 + e- O2 - 1g2

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 partners notebookCompound 1: 76.0% Ru and 24.0% O (by mass), Compound 2: 61.2% Ru and 38.8% O (by mass). (a) What is the empirical formula for Compound 1? (b) What is the empirical formula for Compound 2? Upon determining the melting points of these two compounds, you find that the yellow compound melts at 25 C, while the black powder does not melt up to the maximum temperature of your apparatus, 1200 C. (c) What is the identity of the yellow compound? (d) What is the identity of the black compound? (e) Which compound is molecular? (f) Which compound is ionic?

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 = k1I - EA2, where k is a proportionality constant. (a) How does this definition explain why the electronegativity of F is greater than that of Cl even though Cl 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 F under this definition. (d) Use your result from part (c) to determine the electronegativities of Cl and O using this scale. (e) Another scale for electronegativity defines electronegativity as the average of an atoms first ionization energy and its electron affinity. Using this scale, calculate the electronegativities for the halogens, and scale them so fluorine has an electronegativity of 4.0. On this scale, what is Brs electronegativity?

The compound chloral hydrate, known in detective stories as knockout drops, is composed of 14.52% C, 1.83% H, 64.30% Cl, and 13.35% O by mass, and has a molar mass of 165.4 g>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 Cl atoms bond to a single C atom and that there are a CC bond and two CO bonds in the compound.

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

Acetylene 1C2H22 and nitrogen 1N22 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. Which compound is more stable? (c) Write balanced chemical equations for the complete oxidation of N2 to form N2O51g2 and of acetylene to form CO21g2 and H2O1g2. (d) Calculate the enthalpy of oxidation per mole for N2 and for C2H2 (the enthalpy of formation of N2O51g2 is 11.30 kJ>mol). (e) Both N2 and C2H2 possess triple bonds with quite high bond enthalpies (Table 8.4). Calculate the enthalpy of hydrogenation per mole for both compounds: acetylene plus H2 to make methane, CH4; nitrogen plus H2 to make ammonia, NH3.

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% S and 30.4% N. Measurements of its molecular mass yield a value of 184.3 g>mol. 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 SS distance in the S8 ring is 2.05 .) (d) The enthalpy of formation of the compound is estimated to be 480 kJ>mol -1 . Hf of S(g) is 222.8 kJ>mol -1 . Estimate the average bond enthalpy in the compound.

A common form of elemental phosphorus is the tetrahedral P4 molecule, where all four phosphorus atoms are equivalent: P4 At room temperature phosphorus is a solid. (a) Are there any lone pairs of electrons in the P4 molecule? (b) How many PP bonds are there in the molecule? (c) Draw a Lewis structure for a linear P4 molecule that satisfies the octet rule. Does this molecule have resonance structures? (d) On the basis of formal charges, which is more stable, the linear molecule or the tetrahedral molecule?

Consider benzene 1C6H62 in the gas phase. (a) Write the reaction for breaking all the bonds in C6H61g2, and use data in Appendix C to determine the enthalpy change for this reaction. (b) Write a reaction that corresponds to breaking all the carboncarbon bonds in C6H61g2. (c) By combining your answers to parts (a) and (b) and using the average bond enthalpy for CH from Table 8.4, calculate the average bond enthalpy for the carboncarbon bonds in C6H61g2. (d) Compare your answer from part (c) to the values for CC single bonds and CC double bonds in Table 8.4. Is benzenes CC bond enthalpy exactly halfway between them? If not, which bond type is more similar to that of benzene, CC single or CC double bonds?

Average bond enthalpies are generally defined for gas-phase molecules. Many substances are liquids in their standard state. (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) BrBr, from Br21l2; (b) CCl, from CCl41l2; (c) OO, from H2O21l2 (assume that the OH bond enthalpy is the same as that in the gas phase). (d) Does the process of breaking bonds in the liquid as compared to the gas phase cost more energy? Explain the difference in the H values between the two phases.

Silicon, the element, is the heart of integrated circuits and computer chips in almost all of our electronic devices. Si has the same structure as diamond; each atom is singly bonded to four neighbors. Unlike diamond, silicon has a tendency to oxidize (to SiO2, another extended solid) if exposed to air. (a) Estimate the enthalpy of reaction for the conversion of 1 cm3 of silicon into SiO2. (b) Unlike carbon, silicon rarely forms multiple bonds. Estimate the bond enthalpy of the SiSi bond, assuming that the ratio of the SiSi double bond enthalpy to that of the SiSi single bond is the same as that for carboncarbon bonds.

You have learned that the resonance of benzene, \(\mathrm{C}_{6} \mathrm{H}_{6}\), gives the compound special stability. (a) By using data in Appendix C, compare the heat of combustion of \(1.0 \mathrm{~mol} \mathrm{C}_{6} \mathrm{H}_{6}(\mathrm{~g})\) to the heat of combustion of \(3.0 \mathrm{~mol}\) acetylene, \(\mathrm{C}_{2} \mathrm{H}_{2}(\mathrm{~g})\). Which has the greater fuel value, \(1.0 \mathrm{~mol} \mathrm{C}_{6} \mathrm{H}_{6}(\mathrm{~g}) \text { or } 3.0 \mathrm{~mol} \mathrm{C}_{2} \mathrm{H}_{2}(\mathrm{~g})\)? Are your calculations consistent with benzene being especially stable? (b) Repeat part (a), with the appropriate molecules, for toluene \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CH}_{3}\right)\), a derivative of benzene that has \(a-\mathrm{CH}_{3}\) group in place of one H. (c) Another reaction you can use to compare molecules is hydrogenation, the reaction of a carbon-carbon double bond with \(\mathrm{H}_{2}\) to make a \(\mathrm{C}-\mathrm{C}\) single bond and two \(\mathrm{C}-\mathrm{H}\) single bonds. The experimental heat of hydrogenation of benzene to make cyclohexane \(\mathrm{C}_{6} \mathrm{H}_{12}\), a six-membered ring with 6 \(\mathrm{C}-\mathrm{C}\) single bonds with 12 \(\mathrm{C}-\mathrm{H}\) bonds) is \(208 \mathrm{~kJ} / \mathrm{mol}\). The experimental heat of hydrogenation of cyclohexene \(\mathrm{C}_{6} \mathrm{H}_{10}\), a six-membered ring with one \(\mathrm{C}=\mathrm{C}\) double bond, 5 \(\mathrm{C}-\mathrm{C}\) single bonds, and 10 \(\mathrm{C}-\mathrm{H}\) bonds) to make cyclohexane is \(120 \mathrm{~kJ} / \mathrm{mol}\). Show how these data can provide you with an estimate of the resonance stabilization energy of benzene. (d) Are the bond lengths or angles in benzene, compared to other hydrocarbons, sufficient to decide if benzene exhibits resonance and is especially stable? Discuss. (e) Consider cyclooctatetraene, \(\mathrm{C}_{8} \mathrm{H}_{8}\), which has the octagonal structure shown below. What experiments or calculations could you perform to determine whether cyclooctatetraene exhibits resonance? Equation Transcription: Text Transcription: C_6H_6 1.0 mol C_6H_6(g) 3.0 mol C_2H_2(g) 1.0 mol C_6H_6(g) or 3.0 mol C_2H_2(g) (C_6H_5CH_3) a-CH_3 H_2 C-C C-H C_6H_12 C-C C-H 208 kJ/mol C_6H_10 C=C C-C C-H 120 kJ/mol C_8H_8

Problem 1E Visualizing Concepts For each of these Lewis symbols, indicate the group in the periodic table in which the element X belongs: [Section] (a) X, (b) X, (c) X

Illustrated are four ions \(-A, B, X \text { and } Y-\) showing their relative ionic radii. The ions shown in red carry positive charges: a \(2+\) charge for A and a \(1+\) charge for 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 the combinations in part (a), which leads to the ionic compound having the largest lattice energy? [Section 8.2] Equation Transcription: Text Transcription: -A, B, X and Y- 2+ 1+ 1:1

A portion of a two-dimensional “slab” of \(\mathrm{NaCl}(\mathrm{s})\) is shown here (see Figure 8.3) in which the ions are numbered. The crystal structure of sodium chloride. (a) Which colored balls must represent sodium ions? (b) Which colored balls must represent chloride ions? (c) Consider ion 5. How many attractive electrostatic interactions are shown for it? ( d) Consider ion 5. How many repulsive interactions are shown for it? (e) Is the sum of the attractive interactions in part (c) larger or smaller than the sum of the repulsive interactions in part (d)? (f) If this pattern of ions was extended indefinitely in two dimensions, would the lattice energy be positive or negative? [Section 8.2] Equation Transcription: Text Transcription: NaCl(s)

In the Lewis structure shown here, \(A, D, E, Q, X, \text { and } Z\) represent elements in the first two rows of the periodic table. Identify all six elements so that the formal charges of all atoms are zero. [Section 8.3] Equation Transcription: Text Transcription: A, D, E, Q, X, and Z

The orbital diagram that follows 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] Equation Transcription: Text Transcription: 2+ 4d

Problem 6E Visualizing Concepts Incomplete Lewis structures for the nitrous acid molecule, HNO2, and the nitrite ion, NO2 –, are shown here (a) Complete each Lewis structure by adding electron pairs as needed, (b) Is the formal charge on N the same or different in these two species? (c) Would either HNO2 or NO2– be expected to exhibit resonance? (d) Would you expect the N?O bond in HNO2 to be longer, shorter, or the same length as the N?O bonds in NO2–? [Sections] H?O?N?O O?N?O

The partial Lewis structure that follows 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) How many 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. [Sections 8.3 and 8.8] \(C \stackrel{1}{=} C \stackrel{2}{-} C \stackrel{3}{\equiv} C\) Equation Transcription: Text Transcription: C \stackrel{1}{=} C \stackrel{2}{-} C \stackrel{3}{\equiv} C

Consider the Lewis structure for the polyatomic oxyanion shown here, where X is an element from the third period \((N a-A r)\). By changing the overall charge, \(n\), from \(1-\text { to } 2-\text { to } 3-\) we get three different polyatomic ions. For each of these ions (a) identify the central atom, X; (b) determine the formal charge of the central atom, X; (c) draw a Lewis structure that makes the formal charge on the central atom equal to zero. [Sections 8.5, 8.6, and 8.7] Equation Transcription: Text Transcription: (Na-Ar) n 1- to 2- to 3-

Problem 9E Lewis Symbols (Section) (a) True or false: An element’s number of valence electrons is the same as its atomic number, (b) How many valence electrons does a nitrogen atom possess? (c) An atom has the electron configuration 1s22s22p63s23p2. How many valence electrons does the atom have?

Problem 12E Lewis Symbols (Section) (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 4B. Write the electron configuration for Hf, (c) 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?

Problem 11E Lewis Symbols (Section) Consider the element silicon, Si (a) Write its electron configuration, (b) How many valence electrons does a silicon atom have? (c) Which subshells hold the valence electrons?

Problem 10E Lewis Symbols (Section) (a) True or false: The hydrogen atom is most stable when it has a full octet of electrons, (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 1s22s22p3, how many electrons must it gain to achieve an octet?

Problem 13E Lewis Symbols (Section) Write the Lewis symbol for atoms of each of the following elements: (a) Al, (b) Br, (c) Ar, (d) Sr.

Problem 14E Lewis Symbols (Section) What is the Lewis symbol for each of the following atoms or ions? (a) K, (b) As, (c) Sn2+, (d) N3–

Problem 15E Ionic Bonding (Section) (a) Using Lewis symbols, diagram the reaction between magnesium and oxygen atoms to give the ionic substance MgO, (b) How many electrons are transferred? (c) Which atom loses electrons in the reaction?

Problem 16E Ionic Bonding (Section) (a) Use Lewis symbols to represent the reaction that occurs between Ca and F atoms, (b) What is the chemical formula of the most likely product? (c) How many electrons are transferred? (d) Which atom loses electrons in the reaction?

Problem 17E Ionic Bonding (Section) Predict the chemical formula of the ionic compound formed between the following pairs of elements: (a) Al and F, (b) K and S, (c) Y and O, (d) Mg and N.

Problem 19E Ionic Bonding (Section) Write the electron configuration for each of the following ions, and determine which ones possess noble-gas configurations: (a) Sr2+, (b) Ti2+, (c) Se2–, (d) Ni2+, (e) Br–, (f) Mn3 +

Problem 18E Ionic Bonding (Section) 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) aluminium and oxygen.

Problem 20E Ionic Bonding (Section) Write electron configurations for the following ions, and determine which have noble-gas configurations: (a) Cd2+, (b) P3–, (c) Zr4+, (d) Ru3+, (e) As3–, (f) Ag+

Problem 21E Ionic Bonding (Section) (a) Is lattice energy usually endothermic or exothermic? (b) Write the chemical equation that represents the process of lattice energy for the case of NaCl (c) Would you expect salts like NaCl, which have singly-charged ions, to have larger or smaller lattice energies compared to salts like CaO which are composed of doubly-charged ions?

\(\mathrm{NaCl} and \mathrm{KF}\) 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 \(N a-C l and K-F\) distances. Cation and anion size. Radii, in angstroms, of atoms and their ions for five groups of representative elements. Equation Transcription: Text Transcription: NaCl and KF NaCl and KF Na-Cl or the K-F Na-Cl and K-F

Problem 23E Ionic Bonding (Section) The ionic substances NaF, CaO, and ScN are isoelectronic (they have the same number of electrons). Examine the lattice energies for these substances in Table Make a graph of lattice energy on the vertical axis versus the charge on the cation on the horizontal axis Table Lattice Energies for Some Ionic Compounds Compound Lattice Energy (kJ/mol) Compound Lattice Energy (kJ/mol) LiF 1030 MgCl2 2326 LiCl 834 SrCl2 2127 LiI 730 NaF 910 MgO 3795 NaCl 788 CaO 3414 NaBr 732 SrO 3217 NaI 682 KF 808 ScN 7547 KCl 701 KBr 671 CsCl 657 CsI 600 (a) What is the slope of the line? (b) Make a graph of lattice energy on the vertical axis versus the square of the cation charge on the horizontal axis. What is the slope of this line? (c) Compare how well the data points fall on a line for the graphs in (a) and (b). Which trend is more linear, lattice energy versus cation charge or lattice energy versus cation charge squared? (d) Predict the lattice energy for the compound TiC, if we consider the carbon to have a 4-charge.

Problem 24E Ionic Bonding (Section) (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) Arrange the following substances not listed in Table according to their expected lattice energies, listing them from lowest lattice energy to the highest: MgS, KI, GaN, LiBr. Table Lattice Energies for Some Ionic Compounds Compound Lattice Energy (kJ/mol) Compound Lattice Energy (kJ/mol) LiF 1030 MgCl2 2326 LiCl 834 SrCl2 2127 LiI 730 NaF 910 MgO 3795 NaCl 788 CaO 3414 NaBr 732 SrO 3217 NaI 682 KF 808 ScN 7547 KCl 701 KBr 671 CsCl 657 CsI 600

Consider the ionic compounds \(\mathrm{KF}, \mathrm{NaCl}, \mathrm{NaBr}, \text { and } \mathrm{LiCl}\). (a) Use ionic radii (Figure 7.8) to estimate the cation–anion distance for each compound. Cation and anion size. Radii, in angstroms, of atoms and their ions for five groups of representative elements. (b) Based on your answer to part (a), arrange these four compounds in order of decreasing lattice energy. (c) Check your predictions in part (b) with the experimental values of lattice energy from Table 8.2. Are the predictions from ionic radii correct? Equation Transcription: Text Transcription: KF, NaCl, NaBr, and LiCl

Problem 26E Ionic Bonding (Section) Which of the following trends in lattice energy is due to differences in ionic radii? (a) NaCl >RbBr >CsBr, (b) BaO >KF, (c)SrO >SrCl2.

Problem 28E Ionic Bonding (Section) List the individual steps used in constructing a Born–Haber cycle for the formation of BaI2 from the elements. Which of the steps would you expect to be exothermic?

Problem 27E Ionic Bonding (Section) Energy is required to remove two electrons from Ca to form Ca2+, and energy is required to add two electrons to O to form O2– Yet CaO is stable relative to the free elements. Which statement is the best explanation? (a) The lattice energy of CaO is large enough to overcome these processes, (b) CaO is a covalent compound, and these processes are irrelevant, (c) CaO has a higher molar mass than either Ca or O, (d) The enthalpy of formation of CaO is small, (e) CaO is stable to atmospheric conditions.

(a) Based on the lattice energies of \(\mathrm{MgCl} 2 \text { and } \mathrm{SrCl} 2\) 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 7.11, Figure 7.13 and the value of the second ionization energy for \(\mathrm{Ca}, 1145 \mathrm{~kJ} / \mathrm{mol}\), calculate the lattice energy of \(\mathrm{CaCl} 2\) Equation Transcription: Text Transcription: MgCl2 and SrCl2 CaCl2 Ca, 1145 kJ/mol CaCl2

Use data from Appendix C, Figure 7.10, and Figure 7.12 to calculate the lattice energy of \(R b C l \) Figure 7.12 Electron affinity in \(k J, m o l\) for selected \(s-a n d p-\) block elements. Equation Transcription: Text Transcription: RbCl kJ, mol s- and p-

Problem 31E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) (a) State whether the bonding in each compound is likely to be covalent or not: (i) iron, (ii) sodium chloride, (iii) water, (iv) oxygen, (v) argon (b) A substance XY, formed from two different elements, boils at –33 °C. Is XY likely to be a covalent or an ionic substance?

Problem 32E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) Which of these elements are unlikely to form covalent bonds? S, H, K, Ar, Si.

Problem 33E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) Using Lewis symbols and Lewis structures, diagram the formation of SiCl4 from Si and Cl atoms, showing valence-shell electrons: (a) How many valence electrons does Si have initially? (b) How many valence electrons does each Cl have initially? (c) How many valence electrons surround the Si in the SiCl4 molecule? (d) How many valence electrons surround each Cl in the SiCl4 molecule? (e) How many bonding pairs of electrons are in the SiCl4 molecule?

Problem 34E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) Use Lewis symbols and Lewis structures to diagram the formation of PF3 from P and F atoms, showing valence-shell electrons: (a) How many valence electrons does P have initially? (b) How many valence electrons does each F have initially? (c) How many valence electrons surround the P in the PF3 molecule? (d) How many valence electrons surround each F in the PF3 molecule? (e) How many bonding pairs of electrons are in the PF3 molecule?

Problem 35E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) (a) Construct a Lewis structure for O2 in which each atom achieves an octet of electrons, (b) How many bonding electrons are in the structure? (c) Would you expect the O?O bond in O2 to be shorter or longer than the O?O bond in compounds that contain an O?O single bond? Explain.

Problem 38E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) (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) True or false: The most easily ionizable elements are the most electronegative.

Problem 36E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) (a) Construct a Lewis structure for hydrogen peroxide, H2O2, in which each atom achieves an octet of electrons, (b) How many bonding electrons are between the two oxygen atoms? (c) Do you expect the O?O bond in H2O2 to be longer or shorter than the O?O bond in O2? Explain.

Problem 40E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) By referring only to the periodic table, select (a) the most electronegative element in group 6A, (b) the least electronegative element in the group Al, Si, P, (c) the most electronegative element in the group Ga, P, Cl, Na, (d) the element in the group K, C, Zn, F that is most likely to form an ionic compound with Ba.

Problem 37E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) Which of the following statements about electronegativity is false? (a) Electronegativity is the ability of an atom in a molecule to attract electron density toward itself, (b) Electronegativity is the same thing as electron affinity, (c) The numerical values for electronegativity have no units, (d) Fluorine is the most electronegative element, (e) Cesium is the least electronegative element.

Problem 39E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) Using only the periodic table as your guide, select the most electronegative atom in each of the following sets: (a) Na, Mg, K, Ca, (b) P, S, As, Se, (c) Be, B, C, Si, (d) Zn, Ge, Ga, As.

Problem 42E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) Arrange the bonds in each of the following sets in order of increasing polarity: (a) C?F, O?F, Be?F, (b) O?Cl, S?Br, C?P, (c) C?S, B?F, N?O.

Problem 41E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) Which of the following bonds are polar? (a) B?F, (b) Cl?Cl, (c) Se?O, (d) H?I. Which is the more electronegative atom in each polar bond?

Problem 43E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) (a) From the data in Table calculate the effective charges on the H and Br atoms of the HBr molecule in units of the electronic charge, e. Table Bond Lengths, Electronegativity Differences, and Dipole Moments of the Hydrogen Halides Compound Bond Length (Å) Electro negativity Difference Dipole Moment (D) HF 0.92 1.9 1.82 HCl 1.27 0.9 1.08 HBr 1.41 0.7 0.82 HI 1.61 0.4 0.44 (b) If you were to put HBr under very high pressure, so its bond length decreased significantly, would its dipole moment increase, decrease, or stay the same, if you assume that the effective charges on the atoms do not change?

Problem 44E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) The iodine monobromide molecule, IBr, has a bond length of 2.49 Å and a dipole moment of 1.21 D (a) Which atom of the molecule is expected to have a negative charge? (b) Calculate the effective charges on the I and Br atoms in IBr in units of the electronic charge, e.

Problem 45E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) 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) SiF4 and LaF3, (b) FeCl2 and ReCl6, (c) PbCl4 and RbCl.

Problem 47E Lewis Structures; Resonance Structures (Sections) Draw Lewis structures for the following: (a) SiH4, (b) CO, (c) SF2, (d) H2SO4 (H is bonded to O), (e) ClO2–, (f) NH2OH.

Problem 46E Covalent Bonding, Electronegativity, and Bond Polarity (Sections) 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) TiCl4 and CaF2, (b) ClF3 and VF3, (c) SbCl5 and AlF3.

Problem 49E Lewis Structures; Resonance Structures (Sections) Which one of these statements about formal charge is true? (a) Formal charge is the same as oxidation number, (b) To draw the best Lewis structure, you should minimize formal charge, (c) Formal charge takes into account the different electronegativities of the atoms in a molecule, (d) Formal charge is most useful for ionic compounds, (e) Formal charge is used in calculating the dipole moment of a diatomic molecule.

Problem 48E Lewis Structures; Resonance Structures (Sections) Write Lewis structures for the following: (a) H2CO (both H atoms are bonded to C), (b) H2O2, (c) C2F6 (contains a C?C bond), (d) AsO33–, (e) H2SO3 (H is bonded to O), (f) NH2Cl.

Problem 50E Lewis Structures; Resonance Structures (Sections) (a) Draw the dominant Lewis structure for the phosphorus trifluoride molecule, PF3, (b) Determine the oxidation numbers of the P and F atoms, (c) Determine the formal charges of the P and F atoms.

Problem 51E Lewis Structures; Resonance Structures (Sections) Write Lewis structures that obey the octet rule for each of the following, and assign oxidation numbers and formal charges to each atom: (a) OCS, (b) SOCl2 (S is the central atom), (c) BrO3–, (d) HClO2 (H is bonded to O).

Problem 52E Lewis Structures; Resonance Structures (Sections) 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) SO2, (b) SO3, (c) SO32–, (d) Arrange these molecules/ions in order of increasing S—O bond length.

Problem 53E Lewis Structures; Resonance Structures (Sections) (a) Draw the best Lewis structure(s) for the nitrite ion, NO2–, (b) With what allotrope of oxygen is it isoelectronic? (c) What would you predict for the lengths of the bonds in NO2–relative to N—O single bonds and double bonds?

Problem 55E Lewis Structures; Resonance Structures (Sections) Predict the ordering, from shortest to longest, of the bond lengths in CO, CO2, and CO32–

Problem 54E Lewis Structures; Resonance Structures (Sections) Consider the formate ion, HCO2–, which is the anion formed when formic acid loses an H+ ion. The H and the two O atoms are bonded to the central C atom (a) Draw the best Lewis structure(s) for this ion, (b) Are resonance structures needed to describe the structure? (c) Would you predict that the C—O bond lengths in the formate ion would be longer or shorter relative to those in CO2?

Problem 56E Lewis Structures; Resonance Structures (Sections) Based on Lewis structures, predict the ordering, from shortest to longest, of N—O bond lengths in NO+, NO2 –, and NO3–

Problem 57E Lewis Structures; Resonance Structures (Sections) (a) Do the C—C bond lengths in benzene alternate shortlong-short-long around the ring? Why or why not? (b) Are C—C bond lengths in benzene shorter than C—C single bonds? (c) Are C—C bond lengths in benzene shorter than C?C double bonds?

Mothballs are composed of naphthalene, \(C_{10} H_{8}\), a molecule that consists of two six-membered rings of carbon fused along an edge, as shown in this incomplete Lewis structure: (a) Draw all of the resonance structures of naphthalene. How many are there? (b) Do you expect the \(C-C\) bond lengths in the molecule to be similar to those of \(C-C\) single bonds, \(C=C\) double bonds, or intermediate between \(C-C\) single and \(C=C\) double bonds? (c) Not all of the \(C-C\) bond lengths in naphthalene are equivalent. Based on your resonance structures, how many \(C-C\) bonds in the molecule do you expect to be shorter than the others? Equation Transcription: Text Transcription: C10H8 C-C C-C C=C C-C C=C C-C C-C

Problem 59E Exceptions to the Octet Rule (Section) Indicate whether each statement is true or false: (a) The octet rule is based on the fact that filling in all s and p valence electrons in a shell gives eight electrons, (b) The Si in SiH4 does not follow the octet rule because hydrogen is in an unusual oxidation state, (c) Boron compounds are frequent exceptions to the octet rule because they have too few electrons surrounding the boron. (d) Compounds in which nitrogen is the central atom are frequent exceptions to the octet rule because they have too many electrons surrounding the nitrogen.

Problem 61E Exceptions to the Octet Rule (Section) Draw the dominant Lewis structures for these chlorine–oxygen molecules/ions: ClO, ClO–, ClO2–, ClO3–-, ClO4–-. Which of these do not obey the octet rule?

Problem 62E Exceptions to the Octet Rule (Section) For elements in the third row of the periodic table and beyond, the octet rule is often not obeyed. A friend of yours says this is because these heavier elements are more likely to make double or triple bonds. Another friend of yours says that this is because the heavier elements are larger and can make bonds to more than four atoms at a time. Which friend is most correct?

Problem 60E Exceptions to the Octet Rule (Section) Fill in the blank with the appropriate numbers for both electrons and bonds (considering that single bonds are counted as one, double bonds as two, and triple bonds as three). (a) Fluorine has ____ valence electrons and makes ____ bond(s) in compounds. (b) Oxygen has ____ valence electrons and makes ____ bond(s) in compounds. (c) Nitrogen has ____ valence electrons and makes ____ bond(s) in compounds. (d) Carbon has ____ valence electrons and makes____ bond(s) in compounds.

Problem 63E Exceptions to the Octet Rule (Section) Draw the Lewis structures for each of the following ions or molecules. Identify those in which the octet rule is not obeyed; state which atom in each compound does not follow the octet rule; and state, for those atoms, how many electrons surround these atoms: (a) PH3, (b) AlH3, (c) N3–, (d) CH2Cl2, (e) SnF6.

Problem 64E Exceptions to the Octet Rule (Section) Draw the Lewis structures for each of the following molecules or ions. Identify instances where the octet rule is not obeyed; state which atom in each compound does not follow the octet rule; and state how many electrons surround these atoms: (a) NO, (b) BF3, (c) ICl2–, (d) OPBr3 (the P is the central atom), (e) XeF4.

Problem 65E Exceptions to the Octet Rule (Section) In the vapor phase, BeCl2 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 structures are possible that satisfy the octet rule? (c) On the basis of the formal charges, which Lewis structure is expected to be dominant for BeCl2?

Problem 67E Exceptions to the Octet Rule (Section) Consider the following statement: “For some molecules and ions, a Lewis structure that satisfies the octet rule does not lead to the lowest formal charges, and a Lewis structure that leads to the lowest formal charges does not satisfy the octet rule.” Illustrate this statement using the hydrogen sulfite ion, HSO3–, as an example (the H atom is bonded to one of the O atoms).

Problem 66E Exceptions to the Octet Rule (Section) (a) Describe the molecule xenon trioxide, XeO3, using four possible Lewis structures, one each with zero, one, two, or three Xe?O double bonds, (b) Do any of these resonance structures satisfy the octet rule for every atom in the molecule? (c) Do any of the four Lewis structures have multiple resonance structures? If so, how many resonance structures do you find? (d) Which of the Lewis structures in (a) yields the most favorable formal charges for the molecule?

Problem 68E Exceptions to the Octet Rule (Section) Some chemists believe that satisfaction of the octet rule should be the top criterion for choosing the dominant Lewis structure of a molecule or ion. Other chemists believe that achieving the best formal charges should be the top criterion. Consider the dihydrogen phosphate ion, H2PO4–, in which the H atoms are bonded to O atoms. (a) What is the predicted dominant Lewis structure if satisfying the octet rule is the top criterion? (b) What is the predicted dominant Lewis structure if achieving the best formal charges is the top criterion?

Bond Enthalpies (Section 8.8) Using Table 8.4, estimate ?H for each of the following gas-phase reactions (note that lone pairs on atoms are not shown):

Problem 70E Bond Enthalpies (Section 8.8) Using Table 8.4, estimate ?H for the following gas-phase reactions:

Bond Enthalpies (Section 8.8) Using Table 8.4, estimate ?H for each of the following reactions: (a) 2CH4 (g) + O2 (g) ? 2CH3OH (g) (b) H2 (g) + Br2 (l) ? 2HBr (g) (c) 2H2O2 (g) ? 2H2O (g) + O2 (g)

Bond Enthalpies (Section 8.8) Ammonia is produced directly from nitrogen and hydrogen by using the Haber process, which is perhaps the most widely used industrial chemical reaction on Earth. The chemical reaction is N2 (g) + 3H2 (g) ? 2NH3 (g) a) Use Table 8.4 to estimate the enthalpy change for the reaction. Is it exothermic or endothermic? (b) Calculate the enthalpy change as obtained using ?Hf° values.

Bond Enthalpies (Section) (a) Use bond enthalpies to estimate the enthalpy change for the reaction of hydrogen with ethylene: (b) Calculate the standard enthalpy change for this reaction, using heat of formation.

Bond Enthalpies (Section 8.8) Use Table 8.4 to estimate the enthalpy change for each of the following reactions: (a) H2C ? O (g) + HCl (g) ? H3C ? O ? Cl (g) (b) H2O2 (g) + 2CO (g) ? H2 (g) + 2CO2 (g) (c) 3H2C ? CH2 (g) ? C6H12 (g) (the six carbon atoms form a six-membered ring with two H atoms on each C atom)

Bond Enthalpies (Section) Given the following bond-dissociation energies, calculate the average bond enthalpy for the Ti?Cl bond. ?H(kJ>mol)

Bond Enthalpies (Section) (a) Using average bond enthalpies, predict which of the following reactions will be most exothermic: (i) (ii) (iii) (b) Make a graph that plots the reaction enthalpy you calculated on the vertical axis versus oxidation state of carbon on the horizontal axis. Draw the best-fit line through your data points. Is the slope of this line negative or positive? Does this mean that as the oxidation state of the carbon increases, the reaction with elemental fluorine becomes more or less exothermic? (c) Predict the reaction enthalpy for the reaction of carbonate ion with fluorine from your graph, and compare to what you predict using average bond enthalpies.

Further Applications of the Ideal-Gas Equation (Section) The molar mass of a volatile substance was determined by the Dumas-bulb method described in Exercise 10.53. The unknown vapor had a mass of 0.846 g: the volume of the bulb was , pressure 752 torr. and temperature . Calculate the molar mass of the unknown vapor.

Further Applications of the Ideal-Gas Equation (Section) Magnesium can be used as a 'getter' in evacuated enclosures to react with the last traces of oxygen. (The magnesium is usually heated by passing an electric current through a wire or ribbon of the metal. ) If an enclosure of 0.452 L has a partial pressure of of at 27 SC. What mass of magnesium will react according to the following equation?

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) \(N a^{+}, B r^{-}\) ; (b) \(R b^{+}, B r^{-}\); (c) \(S r^{2+}, S^{2-}\) . Equation Transcription: Text Transcription: N a^{+}, B r^{-} R b^{+}, B r^{-} S r^{2+}, S^{2-}

Further Applications of the Ideal-Gas Equation (Section) The metabolic oxidation of glucose. . in our bodies produces . which is expelled from our lungs as a gas: (a) Calculate the volume of dry C02 produced at body temperature ( ) and 0.970 atm when 24.5 g of glucose is consumed in this reaction (b) Calculate the volume of oxygen you would need, at 1.00 atm and 298 K. to completely oxidize 50.0 g of glucose.

Both Jacques Charles and Joseph Louis Guy-Lussac were avid balloonists. In his original flight in 1783 , Jacques Charles used a balloon that contained approximately of . He generated the using the reaction between iron and hydrochloric acid: How many kilograms of iron were needed to produce this volume of if the temperature was ?

Further Applications of the Ideal-Gas Equation (Section) Hydrogen gas is produced when zinc reacts with sulfuric acid: If 159 mL of wet is collected over water at and a barometric pressure of 738 torr. how many grams of Zn have been consumed (The vapor pressure of water is tabulated in Appendix B.)

Further Applications of the Ideal-Gas Equation (Section) Acetylene gas. . can be prepared by the reaction of calcium carbide with water: Calculate the volume of that is collected over water at by reaction of 1.524 g of if the total pressure of the gas is 753 torr. (The vapor pressure of water is tabulated in Appendix B.)

From the ionic radii given in Figure 7.8, calculate the potential energy of a \(\mathrm{Ca}^{2+} \text { and } \mathrm{O}^{2-}\) ion pair that is 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. Equation Transcription: Text Transcription: Ca^2+ and O^2- CaO

Problem 85AE Consider the collection of nonmetallic elements O, P, Te, I and 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 XY2? (d) Which combinations of elements would likely yield a compound of empirical formula X2Y3?

Problem 86AE The substance chlorine monoxide, ClO(g), is important in atmospheric processes that lead to depletion of the ozone layer. The ClO molecule has an experimental dipole moment of 1.24 D, and the Cl—O bond length is 1.60 Å. (a) Determine the magnitude of the charges on the Cl and O atoms in units of the electronic charge, e. (b) Based on the electronegativities of the elements, which atom would you expect to have a partial negative charge in the ClO molecule? (c) Using formal charges as a guide, propose the dominant Lewis structure for the molecule. (d) The anion ClO-exists. What is the formal charge on the Cl for the best Lewis structure for ClO-?

(a) Using the electronegativities of \(\mathrm{Br} \text { and } \mathrm{Cl}\), estimate the partial charges on the atoms in the \(\mathrm{Br}-\mathrm{Cl}\) molecule. (b) Using these partial charges and the atomic radii given in Figure 7.8, estimate the dipole moment of the molecule. (c) The measured dipole moment of \(\mathrm{BrCl} \text { is } 0.57 \mathrm{D}\). If you assume the bond length in \(\mathrm{BrCl}\) is the sum of the atomic radii, what are the partial charges on the atoms in \(\mathrm{BrCl}\) using the experimental dipole moment? Equation Transcription: Text Transcription: Br and Cl Br - Cl BrCl is 0.57 D BrCl BrCl

A major challenge in implementing the “hydrogen economy” is finding a safe, lightweight, and compact way of storing hydrogen for use as a fuel. The hydrides of light metals are attractive for hydrogen storage because they can store a high weight percentage of hydrogen in a small volume. For example, can release 5.6% of its mass as upon decomposing to NaH(s), Al(s), and . possesses both covalent bonds, which hold polyatomic anions together, and ionic bonds. (a) Write a balanced equation for the decomposition of . (b) Which element in is the most electronegative? Which one is the least electronegative? (c) Based on electronegativity differences, predict the identity of the polyatomic anion. Draw a Lewis structure for this ion. (d) What is the formal charge on hydrogen in the polyatomic ion?

Problem 89AE Although I3-is known, F3-is not. Which statement is the most correct explanation? (a) Iodine is more likely to be electron-deficient. (b) Fluorine is too small to accommodate three nonbonding electron pairs and two bonding electron pairs. (c) Fluorine is too electronegative to form anions. (d) I2 is known but F2 is not. (e) Iodine has a larger electron affinity than fluorine.

(a) Determine the formal charge on the chlorine atom in the hypochlorite ion, , and the perchlorate ion, , using resonance structures where the Cl atom has an octet. (b) What are the oxidation numbers of chlorine in and in ? (c) Perchlorate is a much stronger oxidizing agent than hypochlorite. Suggest an explanation.

(a) Triazine, is like benzene except that in triazine every other C – 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.

Calculate the formal charge on the indicated atom in each of the following molecules or ions: (a) the central oxygen atom in , (b) phosphorus in , (c) nitrogen in , (d) iodine in , (e) chlorine in (hydrogen is bonded to O).

Ortho-dichlorobenzene, \(\mathrm{C}_{6} \mathrm{H}_{4} \mathrm{Cl}_{2}\), is obtained when two of the adjacent hydrogen atoms in benzene are replaced with \(\mathrm{Cl}\) atoms. A skeleton of the molecule is shown here. (a) Complete a Lewis structure for the molecule using bonds and electron pairs as needed. (b) Are there any resonance structures for the molecule? If so, sketch them. (c) Are the resonance structures in (a) and (b) equivalent to one another as they are in benzene? Equation Transcription: Text Transcription: C6H4Cl2 Cl

The following three Lewis structures can be drawn for \(N_{2} O\): (a) Using formal charges, which of these three resonance forms is likely to be the most important? (b) The \(N-N \) bond length in \(N_{2} O \text { is } 1.12 \AA\), slightly longer than a typical \(N=N\) bond; and the \(N-O\) bond length is \(1.19 \AA\), slightly shorter than a typical bond (see Table 8.5). Based on these data, which resonance structure best represents \(N_{2} O\)? Equation Transcription: Å Å Text Transcription: N2O N-N N2O is 1.12Å N=N N-O 1.19Å N2O

Consider the hypothetical molecule B–A ? B. Are the following statements true or false? (a) This molecule cannot exist. (b) If resonance was important, the molecule would have identical A–B bond lengths.

An important reaction for the conversion of natural gas to other useful hydrocarbons is the conversion of methane to ethane. In practice, this reaction is carried out in the presence of oxygen, which converts the hydrogen produced into water. Use Table 8.4 to estimate ?H for these two reactions. Why is the conversion of methane to ethane more favorable when oxygen is used?

Problem 98AE 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 g of nitroglycerin. (b) Write a balanced equation for the decomposition of TNT. Assume that, upon explosion, TNT decomposes into N2(g), CO2(g), H2O(g), and C(s).

Two compounds are isomers if they have the same chemical formula but different arrangements of atoms. Use Table 8.4 to estimate ?H for each of the following gas-phase isomerization reactions and indicate which isomer has the lower enthalpy.

The “plastic” explosive \(C-4\), often used in action movies, contains the molecule cyclotrimethylenetrinitramine, which is often called \(R D X\) (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 \(\mathrm{CO}(\mathrm{g}), \mathrm{N}_{2}(\mathrm{~g}), \text { and } \mathrm{H}_{2} \mathrm{O}(\mathrm{g})\). 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 \(5.0 \mathrm{~g} \text { of } R D X\) decomposes. Equation Transcription: Text Transcription: C-4 RDX C O(g), N2(g), and H2O(g) 5.0 g of RDX

The figure shows the three lowest regions of Earth's atmosphere, (a) Name each and indicate the approximate elevations at which the boundaries occur, (b) In which region is ozone a pollutant7 In which region does it filter UV solar radiation? (c) In which region is infrared radiation from Earth's surface most strongly reflected back? (d) An aurora borealis is due to excitation of atoms and molecules in the atmosphere 55-95 km above Earth's surface. Which regions in the figure are involved in an aurora borealis? (e) Compare the changes in relative concentrations of water vapor and carbon dioxide with increasing elevation in these three regions.

Problem 101IE The Ti2 + ion is isoelectronic with the Ca atom. (a) Write the electron configurations of Ti2 + and Ca. (b) Calculate the number of unpaired electrons for Ca and for Ti2 + . (c) What charge would Ti have to be isoelectronic with Ca2 + ?

(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 \(\mathrm{Sr}(g) \text { is } 1064 \mathrm{~kJ} / \mathrm{mol}\). Use this fact along with data in Appendix C, Figure 7.10, Figure 7.12, and Table 8.2 to calculate \(\Delta \mathrm{H}_{f}^{\circ} \text { for } \mathrm{SrCl}_{2}(s)\). Equation Transcription: Text Transcription: SrCl_2(s) Sr(g) is 1064 kJ/mol \Delta H_f^o for SrCl_2(s)

The electron affinity of oxygen is \(-141 \mathrm{~kJ} / \mathrm{mol}\), corresponding to the reaction \(\mathrm{O}_{2}(\mathrm{~g})+e^{-} \rightarrow O^{-}(\mathrm{g})\) The lattice energy of \(\mathrm{K}_{2} O(\mathrm{~s}) \text { is } 2238 \mathrm{~kJ} / \mathrm{mol}\), Use these data along with data in Appendix C and Figure 7.10 to calculate the “second electron affinity” of oxygen, corresponding to the reaction \(O^{-}(\mathrm{g})+e^{-} \rightarrow O^{2-}(\mathrm{g})\) Equation Transcription: Text Transcription: -141 kJ/mol O2(g)+e-O-(g) K2O(s) is 2238 kJ/mol O-(g)+e-O2-(g)

Problem 105IE 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(I-EA), where k is a proportionality constant. (a) How does this definition explain why the electronegativity of F is greater than that of Cl even though Cl 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, determine the value of k that would lead to an electronegativity of 4.0 for F under this definition. (d) Use your result from part (c) to determine the electronegativities of Cl and O using this scale. (e) Another scale for electronegativity defines electronegativity as the average of an atom’s first ionization energy and its electron affinity. Using this scale, calculate the electronegativities for the halogens, and scale them so fluorine has an electronegativity of 4.0. On this scale, what is Br’s electronegativity?

Problem 104IE 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% O (by mass). (a) What is the empirical formula for Compound 1? (b) What is the empirical formula for Compound 2? Upon determining the melting points of these two compounds, you find that the yellow compound melts at 25 °C, while the black powder does not melt up to the maximum temperature of your apparatus, 1200 °C. (c) What is the identity of the yellow compound? (d) What is the identity of the black compound? (e) Which compound is molecular? (f) Which compound is ionic?

Problem 107IE 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.

Acetylene (C2H2) and nitrogen (N2) 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. Which compound is more stable? (c) Write balanced chemical equations for the complete oxidation of N2 to form N2O5 (g) and of acetylene to form CO2 (g) and H2O (g). (d) Calculate the enthalpy of oxidation per mole for N2 and for C2H2 (the enthalpy of formation of N2O5 (g) is 11.30 kJ/mol). (e) Both N2 and C2H2 possess triple bonds with quite high bond enthalpies (Table 8.4). Calculate the enthalpy of hydrogenation per mole for both compounds: acetylene plus H2 to make methane, CH4; nitrogen plus H2 to make ammonia, NH3.

Problem 109IE 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% S and 30.4% N. Measurements of its molecular mass yield a value of 184.3 g/mol. 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 S–S distance in the S8 ring is 2.05 Å.), (d) The enthalpy of formation of the compound is estimated to be 480 kJ/mol-1. ?Hf° of S(g) is 222.8 kJ/mol-1. Estimate the average bond enthalpy in the compound.

A common form of elemental phosphorus is the tetrahedral \(P_{4}\) molecule, where all four phosphorus atoms are equivalent: At room temperature phosphorus is a solid. (a) Are there any lone pairs of electrons in the \(P_{4}\) molecule? (b) How many \(P-P\) bonds are there in the molecule? (c) Draw a Lewis structure for a linear \(P_{4}\) molecule that satisfies the octet rule. Does this molecule have resonance structures? (d) On the basis of formal charges, which is more stable, the linear molecule or the tetrahedral molecule? Equation Transcription: Text Transcription: P4 P4 P-P P4

Consider benzene (C6H6) in the gas phase. (a) Write the reaction for breaking all the bonds in C6H6 (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 C6H6 (g). (c) By combining your answers to parts (a) and (b) and using the average bond enthalpy for C ? H from Table 8.4, calculate the average bond enthalpy for the carbon–carbon bonds in C6H6 (g). (d) Compare your answer from part (c) to the values for C ? C single bonds and C ? C double bonds in Table 8.4. Is benzene’s C ? C bond enthalpy exactly halfway between them? If not, which bond type is more similar to that of benzene, CC single or CC double bonds?

Problem 106IE The compound chloral hydrate, known in detective stories as knockout drops, is composed of 14.52% C, 1.83% H, 64.30% Cl, and 13.35% O by mass, and has a molar mass of 165.4 g>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 Cl atoms bond to a single C atom and that there are a C – C bond and two C – O bonds in the compound.

Problem 113IE Silicon, the element, is the heart of integrated circuits and computer chips in almost all of our electronic devices. Si has the same structure as diamond; each atom is singly bonded to four neighbors. Unlike diamond, silicon has a tendency to oxidize (to SiO2, another extended solid) if exposed to air. (a) Estimate the enthalpy of reaction for the conversion of 1 cm3 of silicon into SiO2. (b) Unlike carbon, silicon rarely forms multiple bonds. Estimate the bond enthalpy of the Si“Si bond, assuming that the ratio of the Si“Si double bond enthalpy to that of the Si – Si single bond is the same as that for carbon–carbon bonds.

Average bond enthalpies are generally defined for gas-phase molecules. Many substances are liquids in their standard state. (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) Br ? Br, from Br2 (l); (b) C ? Cl, from CCl4 (l); (c) O ? O, from H2O2 (l) (assume that the O ? H bond enthalpy is the same as that in the gas phase). (d) Does the process of breaking bonds in the liquid as compared to the gas phase cost more energy? Explain the difference in the ?H values between the two phases.