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Chemistry The Central Science 12th edition Theodore Brown, Eugene LeMay, Bruce Bursten, Catherine Murphy, Patrick Woodward - Solutions

The ionic substance strontium oxide, SrO, forms from the reaction of strontium metal with molecular oxygen. The arrangement of the ions in solid SrO is analogous to that in solid NaCl (Figure 2.21):(a) Write a balanced equation for the formation of SrO(s) from its elements. (b) Based on the ionic
Explain the variation in ionization energies of carbon, as displayed in this graph:
Group 4A elements have much more negative electron affinities than their neighbors in groups 3A and 5A (see Figure 7.11). Suggest an explanation.
(a) Use orbital diagrams to illustrate what happens when an oxygen atom gains two electrons. (b) Why does O3- not exist?
Use electron configurations to explain the following observations: (a) The first ionization energy of phosphorus is greater than that of sulfur. (b) The electron affinity of nitrogen is lower (less negative) than those of both carbon and oxygen. (c) The second ionization energy of oxygen is
The electron affinities, in kJ / mol, for the group 1B and group 2B metals are(a) Why are the electron affinities of the group 2B elements greater than zero? (b) Why do the electron affinities of the group 1B elements become more negative as we move down the group?
Hydrogen is an unusual element because it behaves in some ways like the alkali metal elements and in other ways like nonmetals. Its properties can be explained in part by its electron configuration and by the values for its ionization energy and electron affinity. (a) Explain why the electron
The first ionization energy of the oxygen molecule is the energy required for the following process: O2 (g) → O2+ (g) + e- The energy needed for this process is 1175kJ/mol, very similar to the first ionization energy of Xe. Would you expect O2 to react with F2? If so, suggest a product or
A certain AB4 molecule has a "seesaw" shape:From which of the fundamental geometries shown in Figure 9.3 could you remove one or more atoms to create a molecule having this seesaw shape?
The diagram that follows shows the highest-energy occupied MOs of a neutral molecule CX, where element X is in the same row of the periodic table as C.(a) Based on the number of electrons, can you determine the identity of X?(b) Would the molecule be diamagnetic or paramagnetic?(c) Consider the
The azide ion, N3-, is linear with two N-N bonds of equal length, 1.16 Ã….(a) Draw a Lewis structure for the azide ion.(b) With reference to Table 8.5, is the observed bond length consistent with your Lewis structure?(c) What hybridization scheme would you expect at each of the nitrogen
In ozone, O3, the two oxygen atoms on the ends of the molecule are equivalent to one another. (a) What is the best choice of hybridization scheme for the atoms of ozone? (b) For one of the resonance forms of ozone, which of the orbitals are used to make bonds and which are used to hold nonbonding
Butadiene, C4H6, is a planar molecule that has the following carbon-carbon bond lengths:(a) Predict the bond angles around each of the carbon atoms and sketch the molecule. (b) Compare the bond lengths to the average bond lengths listed in Table 8.5. Can you explain any differences?
The following sketches show the atomic orbital wave functions (with phases) used to construct some of the MOs of a homo-nuclear diatomic molecule. For each sketch, determine the type of MO that will result from mixing the atomic orbital wave functions as drawn. Use the same labels for the MOs as in
Write the electron configuration for the first excited state for N2 - that is, the state with the highest-energy electron moved to the next available energy level. (a) Is the nitrogen in its first excited state diamagnetic or paramagnetic? (b) Is the N-N bond strength in the first excited state
Azo dyes are organic dyes that are used for many applications, such as the coloring of fabrics. Many azo dyes are derivatives of the organic substance azo-benzene, C12H10N2. A closely related substance is hydrazo-benzene, C12H12N2. The Lewis structures of these two substances are(Recall the
(a) Using only the valence atomic orbitals of a hydrogen atom and a fluorine atom, and following the model of Figure 9.46, how many MOs would you expect for the HF molecule?(b) How many of the MOs from part (a) would be occupied by electrons?(c) It turns out that the difference in energies between
Carbon monoxide, CO, is iso-electronic to N2.(a) Draw a Lewis structure for CO that satisfies the octet rule.(b) Assume that the diagram in Figure 9.46 can be used to describe the MOs of CO. What is the predicted bond order for CO? Is this answer in accord with the Lewis structure you drew in part
The energy-level diagram in Figure 9.36 shows that the sideways overlap of a pair of p orbitals produces two molecular orbitals, one bonding and one anti-bonding. In ethylene there is a pair of electrons in the bonding Ï€ orbital between the two carbons. Absorption of a photon of the appropriate
A compound composed of 2.1% H, 29.8% N, and 68.1% O has a molar mass of approximately 50 g/mol. (a) What is the molecular formula of the compound? (b) What is its Lewis structure if H is bonded to O? (c) What is the geometry of the molecule? (d) What is the hybridization of the orbitals around
An AB2 molecule is described as linear, and the A-B bond length is known. (a) Does this information completely describe the geometry of the molecule? (b) Can you tell how many nonbonding pairs of electrons are around the A atom from this information?
Sulfur tetra-fluoride (SF4) reacts slowly with O2 to form sulfur tetra-fluoride monoxide (OSF4) according to the following unbalanced reaction:SF4 (g) + O2 (g) †’ OSF4 (g)The O atom and the four F atoms in OSF4 are bonded to a central S atom.(a) Balance the equation.(b) Write a Lewis
The phosphorus tri-halides (PX3) show the following variations in the bond angle X-P-X: PF3, 100.3o, PCl3, PBr3, 101.0o PI3, 102.0o. The trend is generally attributed to the change in the electro-negativity of the halogen. (a) Assuming that all electron domains are the same size, what value of the
The molecule 2-butene, C4H8, can undergo a geometric change called cis-trans isomerization:As discussed in the "Chemistry and Life" box on the chemistry of vision, such transformations can be induced by light and are the key to human vision. (a) What is the hybridization at the two central carbon
(a) Compare the bond enthalpies (Table 8.4) of the carbon-carbon single, double, and triple bonds to deduce an average -bond contribution to the enthalpy. What fraction of a single bond does this quantity represent?(b) Make a similar comparison of nitrogen-nitrogen bonds. What do you observe?(c)
Use average bond enthalpies (Table 8.4) to estimate for the atomization of benzene, C6H6: C6H6 (g) → 6C (g) + 6 H (g) Compare the value to that obtained by using ΔHof data given in Appendix C and Hess's law. To what do you attribute the large discrepancy in the two values?
Many compounds of the transition-metal elements contain direct bonds between metal atoms. We will assume that the z-axis is defined as the metal-metal bond axis. (a) Which of the 3d orbitals (Figure 6.23) can be used to make a σ bond between metal atoms? (b) Sketch the σ3d bonding and σ*3d
The organic molecules shown here are derivatives of benzene in which six-member rings are "fused" at the edges of the hexagons.(a) Determine the empirical formula of benzene and of these three compounds. (b) Suppose you are given a sample of one of the compounds. Could combustion analysis be used
Antibonding molecular orbitals can be used to make bonds to other atoms in a molecule. For example, metal atoms can use appropriate d orbitals to overlap with the π*2p orbitals of the carbon monoxide molecule. This is called d-π back bonding. (a) Draw a coordinate axis system in which the y-axis
Methyl isocyanate, CH3NCO, was made infamous in 1984 when an accidental leakage of this compound from a storage tank in Bhopal, India, resulted in the deaths of about 3,800 people and severe and lasting injury to many thousands more. (a) Draw a Lewis structure for methyl isocyanate. (b) Draw a
(a) Methane (CH4) and the perchlorate ion are both described as tetrahedral. What does this indicate about their bond angles? (b) The NH3 molecule is trigonal pyramidal, while BF3 is trigonal planar. Which of these molecules is flat?
How does a trigonal pyramid differ from a tetrahedron so far as molecular geometry is concerned?
Describe the bond angles to be found in each of the following molecular structures: (a) Planar trigonal, (b) Tetrahedral, (c) Octahedral, (d) Linear.
(a) What is meant by the term electron domain?(b) Explain in what way electron domains behave like the balloons in Figure 9.5.Why do they do so?
(a) How does one determine the number of electron domains in a molecule or ion? (b) What is the difference between a bonding electron domain and a nonbonding electron domain?
Would you expect the nonbonding electron-pair domain in NH3 to be greater or less in size than for the corresponding one in PH3? Explain.
In which of these molecules or ions does the presence of nonbonding electron pairs produce an effect on molecular shape, assuming they are all in the gaseous state? (a) SiH4, (b) PF3, (c) HBr, (d) HCN, (e) SO2.
In which of the following molecules can you confidently predict the bond angles about the central atom, and for which would you be a bit uncertain? Explain in each case. (a) H2S, (b) BCl3, (c) CH3I, (d) CBr4, (e) TeBr4.
How many nonbonding electron pairs are there in each of the following molecules: (a) (CH3)2S, (b) HCN, (c) H2C2, (d) CH3F?
Describe the characteristic electron-domain geometry of each of the following numbers of electron domains about a central atom: (a) 3, (b) 4, (c) 5, (d) 6.
What is the difference between the electron-domain geometry and the molecular geometry of a molecule? Use the water molecule as an example in your discussion. Why do we need to make this distinction?
An AB3 molecule is described as having trigonal bipyramidal electron-domain geometry. How many nonbonding domains are on atom A? Explain.
Give the electron-domain and molecular geometries of a molecule that has the following electron domains on its central atom: (a) Four bonding domains and no nonbonding domains, (b) Three bonding domains and two nonbonding domains, (c) Five bonding domains and one nonbonding domain, (d) Four
What are the electron-domain and molecular geometries of a molecule that has the following electron domains on its central atom? (a) Three bonding domains and no nonbonding domains, (b) Three bonding domains and one nonbonding domain, (c) Two bonding domains and two nonbonding domains.
Give the electron-domain and molecular geometries for the following molecules and ions:(a) HCN(b) SO32-(c) SF4(d) PF6-(e) NH3Cl+(f) N3-
Draw the Lewis structure for each of the following molecules or ions, and predict their electron-domain and molecular geometries: (a) AsF3, (b) CH3+, (c) BrF3, (d) ClO3- (e) XeF2, (f) BrO2-.
The figure that follows shows ball-and-stick drawings of three possible shapes of an AF3 molecule.(a) For each shape, give the electron-domain geometry on which the molecular geometry is based.(b) For each shape, how many nonbonding electron domains are there on atom A?(c) Which of the following
For each molecule (a)-(f), indicate how many different electron- domain geometries are consistent with the molecular geometry showna.b. c. d. e. f.
The figure that follows contains ball-and-stick drawings of three possible shapes of an AF4 molecule.(a) For each shape, give the electron-domain geometry on which the molecular geometry is based.(b) For each shape, how many nonbonding electron domains are there on atom A?(c) Which of the following
Give the approximate values for the indicated bond angles in the following molecules:(a)(b) (c) (d)
In which of the following AFn molecules or ions is there more than one F-A-F bond angle: SiF4, PF5, SF4, AsF3?
The three species NH2-, NH3, and NH4+ have H-N-H bond angles of 105o, 107o, and 109o, respectively. Explain this variation in bond angles.
(a) Explain why BrF4- is square planar, whereas BF4- is tetrahedral. (b) How would you expect the H-X-H bond angle to vary in the series H2O, H2S, H2Se? Explain.
(a) Explain why the following ions have different bond angles: ClO2- and NO2-. Predict the bond angle in each case. (b) Explain why the XeF2 molecule is linear.
What is the distinction between a bond dipole and a molecular dipole moment?
Consider a molecule with formula AX3. Supposing the A-X bond is polar, how would you expect the dipole moment of the AX3 molecule to change as the X-A-X bond angle increases from 100o to 120o?
(a) Does SCl2 have a dipole moment? If so, in which direction does the net dipole point? (b) Does BeCl2 have a dipole moment? If so, in which direction does the net dipole point?
The molecule shown here is difluoromethane (CH2F2), which is used as a refrigerant called R-32.(a) Based on the structure, how many electron domains surround the C atom in this molecule?(b) Would the molecule have a nonzero dipole moment?(c) If the molecule is polar, in what direction will the
(a) The PH3 molecule is polar. Does this offer experimental proof that the molecule cannot be planar? Explain. (b) It turns out that ozone, O3, has a small dipole moment. How is this possible, given that all the atoms are the same?
(a) Consider the AF3 molecules in Exercise 9.29.Which of these will have a nonzero dipole moment? Explain.(b) Which of the AF4 molecules in Exercise 9.30 will have a zero dipole moment?In Exercise 29In Exercise 30
(a) What conditions must be met if a molecule with polar bonds is non-polar? (b) What geometries will signify non-polar molecules for AB2, AB3, and AB4 geometries?
Predict whether each of the following molecules is polar or non-polar: (a) IF (b) CS2 (c) SO3 (d) PCl3 (e) SF6 (f) IF5
Predict whether each of the following molecules is polar or non-polar: (a) CCl4 (b) NH3 (c) SF4 (d) XeF4 (e) CH3Br (f) GaH3
Dichloroethylene (C2H2Cl2) has three forms (isomers), each of which is a different substance. (a) Draw Lewis structures of the three isomers, all of which have a carbon-carbon double bond. (b) Which of these isomers has a zero dipole moment? (c) How many isomeric forms can chloroethylene, C2H3Cl,
Dichlorobenzene, C6H4Cl2, exists in three forms (isomers) called ortho, meta, and para:Which of these has a nonzero dipole moment? Explain.
(a) What is meant by the term orbital overlap? (b) Describe what a chemical bond is in terms of electron density between two atoms.
Draw sketches illustrating the overlap between the following orbitals on two atoms: (a) The 2s orbital on each atom, (b) The 2pz orbital on each atom (assume both atoms are on the z-axis), (c) The 2s orbital on one atom and the 2pz orbital on the other atom.
Consider the bonding in an MgH2 molecule. (a) Draw a Lewis structure for the molecule, and predict its molecular geometry. (b) What hybridization scheme is used in MgH2? (c) Sketch one of the two-electron bonds between an Mg hybrid orbital and an H 1s atomic orbital.
The following plot shows the potential energy of two Cl atoms as a function of the distance between them.(a) To what does energy of zero correspond in this diagram?(b) According to the valence-bond model, why does the energy decrease as the Cl atoms move from a large separation to a smaller one?(c)
How would you expect the extent of overlap of the bonding atomic orbitals to vary in the series IF, ICl, IBr, and I2? Explain your answer.
Fill in the blank spaces in the following chart. If the molecule column is blank, find an example that fulfills the conditions of the rest of the row.
Why are there no sp4 or sp5 hybrid orbitals?
(a) Starting with the orbital diagram of a boron atom, describe the steps needed to construct hybrid orbitals appropriate to describe the bonding in BF3. (b)What is the name given to the hybrid orbitals constructed in (a)? (c) Sketch the large lobes of the hybrid orbitals constructed in part (a).
(a) Starting with the orbital diagram of a sulfur atom, describe the steps needed to construct hybrid orbitals appropriate to describe the bonding in SF2. (b)What is the name given to the hybrid orbitals constructed in (a)? (c) Sketch the large lobes of these hybrid orbitals. (d) Would the
Indicate the hybridization of the central atom in (a) BCl3 (b) AlCl4- (c) CS2 (d) GeH4.
What is the hybridization of the central atom in (a) SiCl4 (b) HCN (c) SO3 (d) TeCl2.
Shown here are three pairs of hybrid orbitals, with each set at a characteristic angle. For each pair, determine the type of hybridization, if any that could lead to hybrid orbitals at the specified angle.
(a) Which geometry and central atom hybridization would you expect in the series BH4-, CH4, NH4+? (b) What would you expect for the magnitude and direction of the bond dipoles in this series? (c) Write the formulas for the analogous species of the elements of period 3; would you expect them to have
(a) Draw a picture showing how two p orbitals on two different atoms can be combined to make a sigma bond. (b) Sketch a π bond that is constructed from p orbitals. (c) Which is generally stronger, a σ bond or a π bond? Explain. (d) Can two s orbitals combine to form a π bond? Explain.
In the series SiF4, PF3, and SF2, estimate the F-X-F bond angle in each case and explain your rationale.
(a) If the valence atomic orbitals of an atom are sp hybridized, how many unhybridized p orbitals remain in the valence shell? How many π bonds can the atom form? (b) Imagine that you could hold two atoms that are bonded together, twist them, and not change the bond length. Would it be easier to
(a) Draw Lewis structures for ethane (C2H6), ethylene (C2H4), and acetylene (C2H2). (b) What is the hybridization of the carbon atoms in each molecule? (c) Predict which molecules, if any, are planar. (d) How many σ and π bonds are there in each molecule? (e) Suppose that silicon could form
The nitrogen atoms in N2 participate in multiple bonding, whereas those in hydrazine, N2H4, do not. (a) Draw Lewis structures for both molecules. (b) What is the hybridization of the nitrogen atoms in each molecule? (c) Which molecule has the stronger N-N bond?
Propylene, C3H6, is a gas that is used to form the important polymer called polypropylene. Its Lewis structure is(a) What is the total number of valence electrons in the propylene molecule? (b) How many valence electrons are used to make Ïƒ bonds in the molecule? (c) How many valence
Ethyl acetate, C4H8O2, is a fragrant substance used both as a solvent and as an aroma enhancer. Its Lewis structure is(a) What is the hybridization at each of the carbon atoms of the molecule? (b) What is the total number of valence electrons in ethyl acetate? (c) How many of the valence
Consider the Lewis structure for glycine, the simplest amino acid:(a) What are the approximate bond angles about each of the two carbon atoms, and what are the hybridizations of the orbitals on each of them? (b) What are the hybridizations of the orbitals on the two oxygens and the nitrogen atom,
(a) What is the difference between a localized π bond and a delocalized one? (b) How can you determine whether a molecule or ion will exhibit delocalized π bonding? (c) Is the π bond in NO2- localized or delocalized?
(a) Write a single Lewis structure for SO3, and determine the hybridization at the S atom. (b) Are there other equivalent Lewis structures for the molecule? (c) Would you expect SO3 to exhibit delocalized π bonding? Explain.
Predict the molecular geometry of each of the following molecules:(a).(b). (c).
What hybridization do you expect for the atom indicated in red in each of the following species? (a) CH3CO2- (b) PH4+ (c) AlF3 (d) H2C==CH-CH2+
(a) What is the difference between hybrid orbitals and molecular orbitals? (b) How many electrons can be placed into each MO of a molecule? (c) Can anti-bonding molecular orbitals have electrons in them?
(a) If you combine two atomic orbitals on two different atoms to make a new orbital, is this a hybrid orbital or a molecular orbital? (b) If you combine two atomic orbitals on one atom to make a new orbital, is this a hybrid orbital or a molecular orbital? (c) Does the Pauli Exclusion Principle
Consider the H2+ ion. (a) Sketch the molecular orbitals of the H2+ ion and draw its energy-level diagram. (b) How many electrons are there in the ion? (c) Draw the electron configuration of the ion in terms of its MOs. (d) What is the bond order in H2+? (e) Suppose that the ion is excited by
(a) Sketch the molecular orbitals of the H2- ion and draw its energy-level diagram. (b) Write the electron configuration of the ion in terms of its MOs. (c) Calculate the bond order in H2- (d) Suppose that the ion is excited by light, so that an electron moves from a lower-energy to a
Draw a picture that shows all three 2p orbitals on one atom and all three 2p orbitals on another atom. (a) Imagine the atoms coming close together to bond. How many σ bonds can the two sets of 2p orbitals make with each other? (b) How many π bonds can the two sets of 2p orbitals make with each
(a) What is the probability of finding an electron on the inter-nuclear axis if the electron occupies a π molecular orbital? (b) For a homo-nuclear diatomic molecule, what similarities and differences are there between the π2p MO made from the 2px atomic orbitals and the π2p MO made from the
(a) What are the relationships among bond order, bond length, and bond energy? (b) According to molecular orbital theory, would either Be2 or Be2+ be expected to exist? Explain.
Explain the following: (a) The peroxide ion, O22-, has a longer bond length than the superoxide ion, O2-. (b) The magnetic properties of B2 are consistent with the π2p MOs being lower in energy than the σ2p MO. (c) The O22+ ion has a stronger O-O bond than O2 itself.
In the hydrocarbon(a) What is the hybridization at each carbon atom in the molecule? (b) How many Ïƒ bonds are there in the molecule? (c) How many Ï€ bonds? (d) Identify all the 120o bond angles in the molecule.
(a) What does the term paramagnetic mean?(b) How can one determine experimentally whether a substance is paramagnetic?(c) Which of the following ions would you expect to be paramagnetic: O2+, N22-, Li2+, O22-? For those ions that are paramagnetic, determine the number of unpaired electrons.

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