CHAPTER EIGHT BONDING: GENERAL CONCEPTS. For Review

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1 APTER EIGT BDIG: GEERAL EPTS or Review 1. Electronegativity is the ability of an atom in a molecule to attract electrons to itself. Electronegativity is a bonding term. Electron affinity is the energy change when an electron is added to a substance. Electron affinity deals with isolated atoms in the gas phase. A covalent bond is a sharing of electron pair(s) in a bond between two atoms. An ionic bond is a complete transfer of electrons from one atom to another to form ions. The electrostatic attraction of the oppositely charged ions is the ionic bond. A pure covalent bond is an equal sharing of shared electron pair(s) in a bond. A polar covalent bond is an unequal sharing. Ionic bonds form when there is a large difference in electronegativity between the two atoms bonding together. This usually occurs when a metal with a small electronegativity is bonded to a nonmetal having a large electronegativity. A pure covalent bond forms between atoms having identical or nearly identical eletronegativities. A polar covalent bond forms when there is an intermediate electronegativity difference. In general, nonmetals bond together by forming covalent bonds, either pure covalent or polar covalent. Ionic bonds form due to the strong electrostatic attraction between two oppositely charged ions. ovalent bonds form because the shared electrons in the bond are attracted to two different nuclei, unlike the isolated atoms where electrons are only attracted to one nuclei. The attraction to another nuclei overrides the added electron-electron repulsions. 2. Anions are larger than the neutral atom and cations are smaller than the neutral atom. or anions, the added electrons increase the electron-electron repulsions. To counteract this, the size of the electron cloud increases, placing the electrons further apart from one another. or cations, as electrons are removed, there are fewer electron-electron repulsions and the electron cloud can be pulled closer to the nucleus. Isoelectronic: same number of electrons. Two variables, the number of protons and the number of electrons, determine the size of an ion. Keeping the number of electrons constant, we only have to consider the number of protons to predict trends in size. The ion with the most protons attracts the same number of electrons most strongly resulting in a smaller size. 3. Lattice energy: the change in energy that takes place when separated gaseous ions are packed together to form an ionic solid. The reason ionic compounds form is the extremely favorable lattice energy value (large and negative). Looking at igure 8.11, there are many processes that occur when forming an ionic compound from the elements in their standard state. Most of these processes (if not all) are unfavorable (endothermic). owever, the large, exothermic lattice energy value dominates and the ionic compound forms. 238

2 APTER 8 BDIG: GEERAL EPTS 239 The lattice energy follows oulomb s law (E Q 1 Q 2 /r). Because Mg has ions with +2 and 2 charges, it will have a more favorable lattice energy than a where the charge on the ions are 1 and +1. The reason Mg has +2 and 2 charged ions and not +1 and 1 charged ions is that lattice energy is more favorable as the charges increase. owever, there is a limit to the magnitude of the charges. To form Mg 3+ 3, the ionization energy would be extremely unfavorable for Mg 2+ since an inner core (n = 2) electron is being removed. The same is true for the electron affinity of 2 ; it would be very unfavorable as the added electron goes into the n = 3 level. The lattice energy would certainly be more favorable for Mg 3+ 3, but the unfavorable ionization energy and electron affinity would dominate making Mg 3+ 3 energetically unfavorable overall. In general, ionic compounds want large charges, but only up to the point where valence electrons are removed or added. When we go beyond the valence shell, the energies become very unfavorable. 4. When reactants are converted into products, reactant bonds are broken and product bonds are formed. Thus, for a reaction should be the energy it takes to break the reactant bonds minus the energy released when bonds are formed. Bond energies give good estimates for gas phase reactions, but give poor estimates when solids or liquids are present. This is because bond energy calculations ignore the attractive forces holding solids and liquids together. Gases have the molecules very far apart and they have minimal (assumed zero) attractive forces. This is not true for solids and liquids where the molecules are very close together. Attractive forces in substances are discussed in hapter 10. or an exothermic reaction, stronger bonds are formed in the products as compared to the strength of the bonds broken in the reactants so energy is released. or endothermic reactions, the product bonds are weaker overall and energy must be absorbed. As the number of bonds increase, bond strength increases and bond length decreases. 5. onmetals, which form covalent bonds, have valence electrons in the s and p orbitals. Since there are 4 total s and p orbitals, there is room for only 8 valence electrons (the octet rule). The valence shell for hydrogen is just the 1s orbital. This orbital can hold 2 electrons, so hydrogen follows the duet rule. Drawing Lewis structures is mostly trial and error. The first step is to sum the valence electrons available. ext, attach the bonded atoms with a single bond. This is called the skeletal structure. In general, the atom listed first in a compound is called the central atom; all other atoms listed after the first atom are attached (bonded) to this central atom. If the skeletal structure is something different, we will generally give you hints to determine how the atoms are attached. The final step in drawing Lewis structures is to arrange the remaining electrons around the various atoms to satisfy the octet rule for all atoms (duet role for ). Be and B are the usual examples for molecules that have fewer than 8 electrons. Be 2 and B 3 only have 4 and 6 total valence electrons, respectively; it is impossible to satisfy the octet rule for Be 2 and B 3 because fewer than 8 electrons are present. All row three and heavier nonmetals can have more than 8 electrons around them, but only if they have to. Always satisfy the octet rule when you can; exceptions to the octet rule occur when there are no other options. f the molecules listed in review question 10, Kr 2, I 3, S 4,

3 240 APTER 8 BDIG: GEERAL EPTS Xe 4, P 5, I 5, and S 6 are all examples of central atoms having more than 8 electrons. In all cases, exceptions occur because they have to. The octet rule cannot be satisfied when there is an odd number of valence electrons. There must be an unpaired electron somewhere in the molecule and molecules do not like unpaired electrons. In general, odd electron molecules are very reactive; they react to obtain an even number of valence electrons. 2 is a good example. 2 has 17 valence electrons; when two 2 molecules react, 2 4, which has 34 valence electrons forms. The octet rule can be satisfied for Resonance occurs when more than one valid Lewis structure can be drawn for a particular molecule. A common characteristic of resonance structures is a multiple bond(s) that moves from one position to another. We say the electrons in the multiple bond(s) are delocalized in the molecule. This helps us rationalize why the bonds in a molecule that exhibit resonance are all equivalent in length and strength. Any one of the resonance structures indicates different types of bonds within that molecule. This is not correct, hence none of the individual resonance structures are correct. We think of the actual structure as an average of all the resonance structures; again this helps explain the equivalent bonds within the molecule that experiment tells us we have. 7. ormal charge: a made up charge assigned to an atom in a molecule or polyatomic ion derived from a specific set of rules. The equation to calculate formal charge is: = (number of valence electrons of the free atom) (number of valence electrons assigned to the atom in the molecule) The assigned electrons are all of the lone pair electrons plus one-half of the bonding electrons. ormal charge can be utilized when more than one nonequivalent resonance structure can be drawn for a molecule. The best structure, from a formal charge standpoint, is the structure that has the atoms in the molecule with a formal charge of zero. or organic compounds, carbon has 4 valence electrons and needs 4 more electrons to satisfy the octet rule. arbon does this by forming 4 bonds to other atoms and by having no lone pairs of electrons. Any carbon with 4 bonds and no lone pairs has a formal charge of zero. ydrogen needs just 1 more electron to obtain the e noble gas electron configuration. ydrogen is always attached with a single bond to one other atom. has 5 valence electrons for a formal charge of zero, will form 3 bonds to other atom(s) for 6 electrons, and the remaining 2 electrons are a lone pair on. xygen will have a formal charge of zero when it is attached to other atom(s) with 2 bonds and has 2 lone pairs. The halogens obtain a formal charge of zero by forming 1 bond to another atom as well as having 3 lone pairs. 8. VSEPR = Valence Shell Electron-Pair Repulsion model. The main postulate is that the structure around a given atom is determined principally by minimizing electron-pair repulsion. Electrons don t like each other, so a molecule adopts a geometry to place the electron pairs about a central atom as far apart as possible. The five base geometries and bond angles are:

4 APTER 8 BDIG: GEERAL EPTS 241 umber of bonded atoms plus lone pairs about a central atom Geometry Bond Angle(s) 2 linear trigonal planar tetrahedral trigonal bipyramid 90, octahedral 90 To discuss deviations from the predicted VSEPR bond angles, let us examine 4, 3, and 2. 4 has the true bond angles, but 3 (107.3 ) and 2 (104.5 ) do not. 4 does not have any lone pairs of electrons about the central atom, while 2 and 3 do. These lone pair electrons require more room than bonding electrons, which tends to compress the angles between the bonding pairs. The bond angle for 2 is the smallest because oxygen has two lone pairs on the central atom; the bond angle is compressed more than in 3 where has only one lone pair. So, in general, lone pairs compress the bond angles to a value slightly smaller than predicted by VSEPR. 9. The two general requirements for a polar molecule are: 1. polar bonds 2. a structure such that the bond dipoles of the polar bonds do not cancel. 4, 4 + 4(7) = 32 valence electrons Xe 4, 8 + 4(7) = 36 e Xe tetrahedral, square planar, 90 S 4, 6 + 4(7) = 34 e S see-saw, 90, 120 The arrows indicate the individual bond dipoles in the three molecules (the arrows point to the more electronegative atom in the bond which will be the partial negative end of the bond dipole). All three of these molecules have polar bonds. To determine the polarity of the

5 242 APTER 8 BDIG: GEERAL EPTS overall molecule, we sum the effect of all of the individual bond dipoles. In 4, the fluorines are symmetrically arranged about the central carbon atom. The net result is for all of the individual bond dipoles to cancel each other out giving a nonpolar molecule. In Xe 4, the 4 Xe bond dipoles are also symmetrically arranged and Xe 4 is also nonpolar. The individual bond dipoles cancel out when summed together. In S 4, we also have 4 polar bonds. But in S 4, the bond dipoles are not symmetrically arranged and they do not cancel each other out. S 4 is polar. It is the positioning of the lone pair that disrupts the symmetry in S 4. 2, 4 + 2(6) = 16 e S, = 16 e S 2 is nonpolar because the individual bond dipoles cancel each other out, but S is polar. By replacing an with a less electronegative S atom, the molecule is not symmetric any more. The individual bond dipoles do not cancel since the S bond dipole is smaller than the bond dipole resulting in a polar molecule. 10. To predict polarity, draw in the individual bond dipoles, then sum up the net effect of the bond dipoles on each other. If the net effect is to have the bond dipoles cancel each other out, then the molecule is nonpolar. If the net effect of the bond dipoles is to not cancel each other out, then the molecule will have a partial positive end and a partial negative end (the molecule is polar). This is called a dipole moment or a polar molecule. 2, 4 + 2(6) = 16 valence electrons S 2, 6 + 2(6) = 18 e S S linear, 180, nonpolar V-shaped, 120, polar Kr 2, 8 + 2(7) = 22 e Kr S 3, 6 + 3(6) = 24 e S + 2 other resonance structures linear, 180, nonpolar trigonal planar, 120, nonpolar

6 APTER 8 BDIG: GEERAL EPTS 243 3, 5 + 3(7) = 26 e trigonal pyramid, < 109.5, polar The bond angles will be somewhat less than due to the lone pair on the central nitrogen atom needing more space. 4, 4 + 4(7) = 32 e tetrahedral, 109.5, nonpolar Xe 4, 8 + 4(7) = 36 e Xe I 3, 7 + 3(7) = 28 e I T-shaped, 90, polar S 4, 6 + 4(7) = 34 e S see-saw, 90 and 120, polar P 5, 5 + 5(7) = 40 e P square planar, 90, nonpolar I 5, 7 + 5(7) = 42 e I trigonal bipyramid, 90 and 120, nonpolar S 6, 6 + 6(7) = 48 e S square pyramid, 90, polar octahedral, 90, nonpolar

7 244 APTER 8 BDIG: GEERAL EPTS Questions 13. f the compounds listed, P 2 5 is the only compound containing only covalent bonds. ( 4 ) 2 S 4, a 3 (P 4 ) 2, K 2, and K are all compounds composed of ions so they exhibit ionic bonding. The ions in ( 4 ) 2 S 4 are 4 + and S 4 2. ovalent bonds exist between the and atoms in 4 + and between the S and atoms in S 4 2. Therefore, ( 4 ) 2 S 4 contains both ionic and covalent bonds. The same is true for a 3 (P 4 ) 2. The bonding is ionic between the a 2+ and P 4 3 ions and covalent between the P and atoms in P 4 3. Therefore, ( 4 ) 2 S 4 and a 3 (P 4 ) 2 are the compounds with both ionic and covalent bonds. 14. Ionic solids are held together by strong electrostatic forces which are omnidirectional. i. or electrical conductivity, charged species must be free to move. In ionic solids the charged ions are held rigidly in place. nce the forces are disrupted (melting or dissolution), the ions can move about (conduct). ii. Melting and boiling disrupts the attractions of the ions for each other. Because these electrostatic forces are strong, it will take a lot of energy (high temperature) to accomplish this. iii. If we try to bend a piece of material, the ions must slide across each other. or an ionic solid the following might happen: strong attraction strong repulsion Just as the layers begin to slide, there will be very strong repulsions causing the solid to snap across a fairly clean plane. iv. Polar molecules are attracted to ions and can break up the lattice. These properties and their correlation to chemical forces will be discussed in detail in hapters 10 and Electronegativity increases left to right across the periodic table and decreases from top to bottom. ydrogen has an electronegativity value between B and in the second row, and identical to P in the third row. Going further down the periodic table, has an electronegativity value between As and Se (row 4) and identical to Te (row 5). It is important to know where hydrogen fits into the electronegativity trend, especially for rows 2 and 3. If you know where fits into the trend, then you can predict bond dipole directions for nonmetals bonded to hydrogen. 16. linear structure (180 bond angle) S polar, bond dipoles do not cancel nonpolar, bond dipoles cancel

8 APTER 8 BDIG: GEERAL EPTS 245 trigonal planar structure (120 bond angle) S + 2 other resonance structures polar, bond dipoles do not cancel nonpolar, bond dipoles cancel tetrahedral structure (109.5 bond angles) polar, bond dipoles do not cancel nonpolar, bond dipoles cancel 17. or ions, concentrate on the number of protons and the number of electrons present. The species whose nucleus holds the electrons most tightly will be smallest. or example, anions are larger than the neutral atom. The anion has more electrons held by the same number of protons in the nucleus. These electrons will not be held as tightly, resulting in a bigger size for the anion as compared to the neutral atom. or isoelectronic ions, the same number of electrons are held by different numbers of protons in the various ions. The ion with the most protons holds the electrons tightest and is smallest in size. 18. Two other factors that must be considered are the ionization energy needed to produce more positively charged ions and the electron affinity needed to produce more negatively charged ions. The favorable lattice energy more than compensates for the unfavorable ionization energy of the metal and for the unfavorable electron affinity of the nonmetal, as long as electrons are added to or removed from the valence shell. nce the valence shell is full, the ionization energy required to remove another electron is extremely unfavorable; the same is true for electron affinity when an electron is added to a higher n shell. These two quantities are so unfavorable after the valence shell is complete, that they overshadow the favorable lattice energy and the higher charged ionic compounds do not form. 19. ossil fuels contain a lot of carbon and hydrogen atoms. ombustion of fossil fuels (reaction with 2 ) produces 2 and 2. Both these compounds have very strong bonds. Because strong bonds are formed, combustion reactions are very exothermic. 20. Statements a and c are true. or statement a, Xe 2 has 22 valence electrons and it is impossible to satisfy the octet rule for all atoms with this number of electrons. The best Lewis structure is: Xe

9 246 APTER 8 BDIG: GEERAL EPTS or statement c, + has 10 valence electrons, while has 12 valence electrons. The Lewis structures are: + Because a triple bond is stronger than a double bond, + has a stronger bond. or statement b, S 4 has 5 electron pairs around the sulfur in the best Lewis structure; it is an exception to the octet rule. Because 4 has the same number of valence electrons as S 4, 4 would also have to be an exception to the octet rule. owever, Row 2 elements like never have more than 8 electrons around them, so 4 does not exist. or statement d, two resonance structures can be drawn for ozone: When resonance structures can be drawn, the actual bond lengths and strengths are all equal to each other. Even though each Lewis structure implies the two bonds are different, this is not the case in real life. In real life, both of the bonds are equivalent. When resonance structures can be drawn, you can think of the bonding as an average of all of the resonance structures , 4 + 2(6) = 16 valence electrons The formal charges are shown above the atoms in the three Lewis structures. The best Lewis structure for 2 from a formal charge standpoint is the first structure having each oxygen double bonded to carbon. This structure has a formal charge of zero on all atoms (which is preferred). The other two resonance structures have nonzero formal charges on the oxygens making them less reasonable. or 2, we usually ignore the last two resonance structures and think of the first structure as the true Lewis structure for nly statement c is true. The bond dipoles in 4 and Kr 4 are arranged in a manner that they all cancel each other out, making them nonpolar molecules ( 4 has a tetrahedral molecular structure while Kr 4 has a square planar molecular structure). In Se 4, the bond dipoles in this see-saw molecule do not cancel each other out, so Se 4 is polar. or statement a, all the molecules have either a trigonal planar geometry or a trigonal bipyramid geometry; both of which have 120 bond angles. owever, Xe 2 has three lone pairs and two bonded fluorine atoms around it. Xe 2 has a linear molecular structure with a 180 bond angle. With three lone pairs, we no longer have a 120 bond angle in Xe 2. or statement b, S 2 has a V- shaped molecular structure with a bond angle of about 120. S 2 is linear with a 180 bond angle and S 2 is V-shaped but with a bond angle. The three compounds do not have the same bond angle. or statement d, central atoms adopt a geometry to minimize electron repulsions, not maximize them.

10 APTER 8 BDIG: GEERAL EPTS 247 Exercises hemical Bonds and Electronegativity 23. The general trend for electronegativity is: 1) increase as we go from left to right across a period and 2) decrease as we go down a group Using these trends, the expected orders are: a. < < b. Se < S < c. Sn < Ge < Si d. Tl < Ge < S 24. a. Rb < K < a b. Ga < B < c. Br < < d. S < < 25. The most polar bond will have the greatest difference in electronegativity between the two atoms. rom positions in the periodic table, we would predict: a. Ge b. P c. S d. Ti 26. a. Sn b. Tl Br c. Si d. 27. The general trends in electronegativity used in Exercises 8.23 and 8.25 are only rules of thumb. In this exercise, we use experimental values of electronegativities and can begin to see several exceptions. The order of E from igure 8.3 is: a. (2.5) < (3.0) < (3.5) same as predicted b. Se (2.4) < S (2.5) < (3.0) same c. Si = Ge = Sn (1.8) different d. Tl (1.8) = Ge (1.8) < S (2.5) different Most polar bonds using actual E values: a. Si and Ge have equal polarity (Ge predicted). b. P (same as predicted) c. S (same as predicted) d. Ti (same as predicted) 28. The order of E from igure 8.3 is: a. Rb (0.8) = K (0.8) < a (0.9), different b. Ga (1.6) < B (2.0) < (3.5), same c. Br (2.8) < (3.0) < (4.0), same d. S (2.5) < (3.5) < (4.0), same

11 248 APTER 8 BDIG: GEERAL EPTS Most polar bonds using actual E values: a. most polar (Sn predicted) b. Al Br most polar (Tl Br predicted). c. Si (same as predicted). d. Each bond has the same polarity, but the bond dipoles point in opposite directions. xygen is the positive end in the bond dipole, and oxygen is the negative end in the bond dipole. ( predicted.) 29. Use the electronegativity trend to predict the partial negative end and the partial positive end of the bond dipole (if there is one). To do this, you need to remember that has electronegativity between B and and identical to P. Answers b, d, and e are incorrect. or d (Br 2 ), the bond between two Br atoms will be a pure covalent bond where there is equal sharing of the bonding electrons and no dipole moment. or b and e, the bond polarities are reversed. In I, the more electronegative atom will be the partial negative end of the bond dipole with I having the partial positive end. In P, the more electronegative oxygen will be the partial negative end of the bond dipole with P having the partial positive end. In the following, we used arrows to indicate the bond dipole. The arrow always points to the partial negative end of a bond dipole (which always is the most electronegative atom in the bond). I P 30. See Exercise 8.29 for a discussion on bond dipoles. We will use arrows to indicate the bond dipoles. The arrow always points to the partial negative end of the bond dipole which will always be to the more electronegative atom. The tail of the arrow indicates the partial positive end of the bond dipole. a. b. P is a pure covalent (nonpolar) bond since P and have identical electronegativities. c. d. Br Te e. The actual electronegativity difference between Se and S is so small that Se S this bond is probably best characterized as a pure covalent bond having no bond dipole. 31. Electronegativity values increase from left to right across the periodic table. The order of electronegativities for the atoms from smallest to largest electronegativity will be = P < < < <. The most polar bond will be since it will have the largest difference in electronegativities, and the least polar bond will be P since it will have the smallest difference in electronegativities (ΔE = 0). The order of the bonds in decreasing polarity will be > > > > P. 32. Ionic character is proportional to the difference in electronegativity values between the two elements forming the bond. Using the trend in electronegativity, the order will be:

12 APTER 8 BDIG: GEERAL EPTS 249 Br Br < < < a < K least most ionic character ionic character ote that Br Br, and bonds are all covalent bonds since the elements are all nonmetals. The a and K bonds are ionic as is generally the case when a metal forms a bond with a nonmetal. Ions and Ionic ompounds 33. r + : [Xe]6s 2 4f 14 5d 10 6p 6 = [Rn]; Be 2+ : 1s 2 = [e]; P 3 : [e]3s 2 3p 6 = [Ar] : [e]3s 2 3p 6 = [Ar]; Se 2 : [Ar]4s 2 3d 10 4p 6 = [Kr] 34. a. Mg 2+ : 1s 2 2s 2 2p 6 ; K + : 1s 2 2s 2 2p 6 3s 2 3p 6 ; Al 3+ : 1s 2 2s 2 2p 6 b. 3, 2 and : 1s 2 2s 2 2p 6 ; Te 2- : [Kr]5s 2 4d 10 5p a. Sc 3+ : [Ar] b. Te 2 : [Xe] c. e 4+ : [Xe] and Ti 4+ : [Ar] d. Ba 2+ : [Xe] All of these ions have the noble gas electron configuration shown in brackets. 36. a. s 2 S is composed of s + and S 2. s + has the same electron configuration as Xe, and S 2 has the same configuration as Ar. b. Sr 2 ; Sr 2+ has the Kr electron configuration and has the e configuration. c. a 3 2 ; a 2+ has the Ar electron configuration and 3 has the e configuration. d. AlBr 3 ; Al 3+ has the e electron configuration and Br has the Kr configuration. 37. There are many possible ions with 54 electrons. Some are: Sb 3, Te 2, I, s +, Ba 2+ and La 3+. In terms of size, the ion with the most protons will hold the electrons the tightest and will be the smallest. The largest ion will be the ion with the fewest protons. The size trend is: La 3+ < Ba 2+ < s + < I < Te 2 < Sb 3 smallest largest 38. All of these ions have 18 e ; the smallest ion (Sc 3+ ) has the most protons attracting the 18 e and the largest ion has the fewest protons (S 2 ). The order in terms of increasing size is: Sc 3+ < a 2+ < K + < < S 2. In terms of the atom size indicated in the question: K + a 2+ Sc 3+ S 2- -

13 250 APTER 8 BDIG: GEERAL EPTS 39. a. u > u + > u 2+ b. Pt 2+ > Pd 2+ > i 2+ c. 2 > > d. La 3+ > Eu 3+ > Gd 3+ > Yb 3+ e. Te 2 > I > s + > Ba 2+ > La 3+ or answer a, as electrons are removed from an atom, size decreases. Answers b and d follow the radii trend. or answer c, as electrons are added to an atom, size increases. Answer e follows the trend for an isoelectronic series, i.e., the smallest ion has the most protons. 40. a. V > V 2+ > V 3+ > V 5+ b. s + > Rb + > K + > a + c. Te 2 > I > s + > Ba 2+ d. P 3 > P 2 > P > P e. Te 2 > Se 2 > S 2 > a. Al 3+ and S 2 are the expected ions. The formula of the compound would be Al 2 S 3 (aluminum sulfide). b. K + and 3 ; K 3, potassium nitride c. Mg 2+ and ; Mg 2, magnesium chloride d. s + and Br ; sbr, cesium bromide 42. a. Ga 3+ and I ; GaI 3, gallium iodide b. a + and 2 ; a 2, sodium oxide or a + and 2 2 ; a 2 2, sodium peroxide c. Sr 2+ and ; Sr 2, strontium fluoride d. a 2+ and P 3 ; a 3 P 2, calcium phosphide 43. Lattice energy is proportional to Q 1 Q 2 /r where Q is the charge of the ions and r is the distance between the ions. In general, charge effects on lattice energy are much greater than size effects. a. a; a + is smaller than K +. b. Li; is smaller than. c. Mg; 2 has a greater charge than -. d. e() 3 ; e 3+ has a greater charge than e 2+. e. a 2 ; 2 has a greater charge than. f. Mg; The ions are smaller in Mg. 44. a. Li; Li + is smaller than s +. b. abr; Br - is smaller than I. c. Ba; 2 has a greater charge than -. d. as 4 ; a 2+ has a greater charge than a +. e. K 2 ; 2 has a greater charge than -. f. Li 2 ; The ions are smaller in Li K(s) K(g) Δ = 64 kj (sublimation) K(g) K + (g) + e Δ = 419 kj (ionization energy) 1/2 2 (g) (g) Δ = 239/2 kj (bond energy) (g) + e (g) Δ = 349 kj (electron affinity) K + (g) + (g) K(s) Δ = 690. kj (lattice energy) K(s) + 1/2 2 (g) K(s) = 437 kj/mol Δ f

14 APTER 8 BDIG: GEERAL EPTS Mg(s) Mg(g) Δ = 150. kj (sublimation) Mg(g) Mg + (g) + e - Δ = 735 kj (IE 1 ) Mg + (g) Mg 2+ (g) + e Δ = 1445 kj (IE 2 ) 2 (g) 2 (g) Δ = 154 kj (BE 2 (g) + 2 e (g) Δ = 2(-328) kj (EA) Mg 2+ (g) (g) Mg 2 (s) Δ = 3916 kj (LE) Mg(s) + 2 (g) Mg 2 (s) = 2088 kj/mol Δ f 47. rom the data given, it takes less energy to produce Mg + (g) + (g) than to produce Mg 2+ (g) + 2 (g). owever, the lattice energy for Mg 2+ 2 will be much more exothermic than that for Mg + due to the greater charges in Mg The favorable lattice energy term dominates and Mg 2+ 2 forms. 48. a(g) a + (g) + e Δ = IE 1 = 495 kj (Table 7.5) (g) + e (g) Δ = EA = kj (Table 7.7) a(g) + (g) a + (g) + (g) Δ = 167 kj The described process is endothermic. What we haven t accounted for is the extremely favorable lattice energy. ere, the lattice energy is a large negative (exothermic) value, making the overall formation of a a favorable exothermic process. 49. Use igure 8.11 as a template for this problem. Li(s) Li(g) Δ sub =? Li(g) Li + (g) + e Δ = 520. kj 1/2 I 2 (g) I(g) Δ = 151/2 kj I(g) + e I (g) Δ = 295 kj Li + (g) + I (g) LiI(s) Δ = 753 kj Li(s) + 1/2 I 2 (g) LiI(s) Δ = 272 kj Δ sub / = 272, Δ sub = 181 kj 50. Let us look at the complete cycle for a 2 S. 2 a(s) 2 a(g) 2 Δ sub, a = 2(109) kj 2 a(g) 2 a + (g) + 2 e 2 IE = 2(495) kj S(s) S(g) S(g) + e S (g) Δ sub, S = 277 kj EA 1 = 200. kj S - (g) + e S 2 (g) EA 2 =? 2 a + (g) + S 2 (g) a 2 S LE = 2203 kj 2 a(s) + S(s) a 2 S(s) Δ f = 365 kj

15 252 APTER 8 BDIG: GEERAL EPTS Δ f = 2 Δ + 2 IE + Δ sub, S + EA 1 + EA 2 + LE, 365 = EA 2, EA 2 = 553 kj sub, a or each salt: Δ f = 2 Δ sub, M + 2 IE LE + EA 2 K 2 S: 381 = 2(90.) + 2(419) EA 2, EA 2 = 576 kj Rb 2 S: 361 = 2(82) + 2(409) EA 2, EA 2 = 529 kj s 2 S: 360. = 2(78) + 2(382) EA 2, EA 2 = 493 kj We get values from 493 to 576 kj. The mean value is: = 538 kj 4 We can represent the results as EA 2 = 540 ± 50 kj. 51. a 2+ has a greater charge than a +, and Se 2 is smaller than Te 2. The effect of charge on the lattice energy is greater than the effect of size. We expect the trend from most exothermic to least exothermic to be: ase > ate > a 2 Se > a 2 Te ( 2862) ( 2721) ( 2130) ( 2095) This is what we observe. 52. Lattice energy is proportional to the charge of the cation times the charge of the anion, Q 1 Q 2. ompound Q 1 Q 2 Lattice Energy e 2 (+2)( 1) = kj/mol e 3 (+3)( 1) = kj/mol e 2 3 (+3)( 2) = 6 14,744 kj/mol Bond Energies 53. a. + 2 Bonds broken: Bonds formed: 1 (432 kj/mol) 2 (427 kj/mol) 1 (239 kj/mol) Δ = ΣD broken ΣD formed, Δ = 432 kj kj 2(427) kj = 183 kj b

16 APTER 8 BDIG: GEERAL EPTS 253 Bonds broken: Bonds formed: 1 (941 kj/mol) 6 (391 kj/mol) 3 (432 kj/mol) Δ = 941 kj + 3(432) kj 6(391) kj = 109 kj 54. Sometimes some of the bonds remain the same between reactants and products. To save time, only break and form bonds that are involved in the reaction. a. + 2 Bonds broken: Bonds formed: 1 (891 kj/mol) 1 (305 kj/mol) 2 (432 kj/mol) 2 (413 kj/mol) 2 (391 kj/mol) Δ = 891 kj + 2(432 kj) - [305 kj + 2(413 kj) + 2(391 kj)] = -158 kj b Bonds broken: Bonds formed: 1 (160. kj/mol) 4 (565 kj/mol) 4 (391 kj/mol) 1 (941 kj/mol) 2 (154 kj/mol) Δ = 160. kj + 4(391 kj) + 2(154 kj) [4(565 kj) kj] = 1169 kj 55. Bonds broken: 1 (305 kj/mol) Bonds formed: 1 (347 kj/mol) Δ = ΣD broken ΣD formed, Δ = = 42 kj

17 254 APTER 8 BDIG: GEERAL EPTS ote: Sometimes some of the bonds remain the same between reactants and products. To save time, only break and form bonds that are involved in the reaction Bonds broken: Bonds formed: 1 (1072 kj/mol) 1 (347 kj/mol) 1 (358 kj/mol) 1 = (745 kj/mol) 1 (358 kj/mol) Δ = [ ] = 20. kj Bonds broken: Bonds formed: 5 (413 kj/mol) 2 2 = (799 kj/mol) 1 (347 kj/mol) 3 2 (467 kj/mol) 1 (358 kj/mol) 1 (467 kj/mol) 3 = (495 kj/mol) Δ = 5(413 kj) kj kj kj + 3(495 kj) [4(799 kj) + 6(467 kj)] /2 = 2 == + = 1276 kj Bonds broken: Bonds formed: 2 (413 kj/mol) 2 2 = (799 kj/mol) 1 (839 kj/mol) 2 (467 kj/mol) 5/2 = (495 kj/mol) = 2(413 kj) kj + 5/2 (495 kj) [4(799 kj) + 2(467 kj] = 1228 kj

18 APTER 8 BDIG: GEERAL EPTS 255 Bonds broken: Bonds formed: 3 (467 kj/mol) 1 = (745 kj/mol); This is not 2. 1 (146 kj/mol) 2 2 (467 kj/mol) 1 (413 kj/mol) 1 (358 kj/mol) = 3(467 kj) kj kj kj [745 kj + 4(467 kj)] = 295 kj Bonds broken: Bonds formed: 9 (160. kj/mol) 24 (467 kj/mol) 4 (305 kj/mol) 9 (941 kj/mol) 12 (413 kj/mol) 8 = (799 kj/mol) 12 (391 kj/mol) 10 = (607 kj/mol) 10 (201 kj/mol) Δ = 9(160.) + 4(305) + 12(413) + 12(391) + 10(607) + 10(201) Δ = 20,388 kj 26,069 kj = kj [24(467) + 9(941) + 8(799)] 61. Bonds broken: Bonds formed: 1 = (614 kj/mol) 1 (347 kj/mol) 1 (154 kj/mol) 2 (D ) + = -549 kj Δ = 549 kj = 614 kj kj [347 kj + 2 D ], 2 D = 970., D = 485 kj/mol 62. Let x = bond energy for A 2, then 2x = bond energy for AB. = 285 kj = x kj [2(2x)], 3x = 717, x = 239 kj/mol

19 256 APTER 8 BDIG: GEERAL EPTS The bond energy for A 2 is 239 kj/mol. 63. a. Δ = 2 f = 2 mol ( 92 kj/mol) = 184 kj (= 183 kj from bond energies), b. Δ = 2 f, 3 = 2 mol ( 46 kj/mol) = 92 kj (= 109 kj from bond energies) omparing the values for each reaction, bond energies seem to give a reasonably good estimate of the enthalpy change for a reaction. The estimate is especially good for gas phase reactions (g) + (g) 3 (l) Δ = 484 kj [(-201 kj) + ( kj)] = 173 kj Using bond energies, Δ = -20. kj. or this reaction, bond energies give a much poorer estimate for Δ as compared to the gas phase reactions in Exercise The reason is that not all species are gases in Exercise Bond energies do not account for the energy changes that occur when liquids and solids form instead of gases. These energy changes are due to intermolecular forces and will be discussed in hapter a. Using S 4 data: S 4 (g) S(g) + 4 (g) Δ = 4 D S = (79.0) ( 775) = kj D S = kj = kj/mol 4 mols bonds Using S 6 data: S 6 (g) S(g) + 6 (g) Δ = 6 D S = (79.0) ( 1209) = 1962 kj kj D S = = kj/mol 6 mol b. The S bond energy in the table is 327 kj/mol. The value in the table was based on the S bond in S 6. c. S(g) and (g) are not the most stable forms of the elements at 25. The most stable forms are S 8 (s) and 2 (g); f = 0 for these two species (g) (g) + 3 (g); Δ = 3 D = (216.0) ( 46.1) = kj D = kj = kj/mol 389 kj/mol 3 mol bonds D calc = 389 kj/mol as compared to 391 kj/mol in the table. There is good agreement.

20 APTER 8 BDIG: GEERAL EPTS 257 Lewis Structures and Resonance 67. Drawing Lewis structures is mostly trial and error. owever, the first two steps are always the same. These steps are 1) count the valence electrons available in the molecule/ion, and 2) attach all atoms to each other with single bonds (called the skeletal structure). Unless noted otherwise, the atom listed first is assumed to be the atom in the middle (called the central atom) and all other atoms in the formula are attached to this atom. The most notable exceptions to the rule are formulas which begin with, e.g., 2, 2, etc. ydrogen can never be a central atom since this would require to have more than two electrons. In these compounds, the atom listed second is assumed to be the central atom. After counting valence electrons and drawing the skeletal structure, the rest is trial and error. We place the remaining electrons around the various atoms in an attempt to satisfy the octet rule (or duet rule for ). Keep in mind that practice makes perfect. After practicing you can (and will) become very adept at drawing Lewis structures. a. has = 10 valence b. P 3 has 5 + 3(1) = 8 valence electrons. P P Skeletal structure Lewis structure Skeletal structure Lewis structure c. 3 has (7) = 26 valence d. + 4 has 5 + 4(1) 1 = 8 valence electrons. electrons. + ote: Subtract valence electrons for positive charged ions. Skeletal structure Lewis structure Lewis structure e. 2 has 2(1) = 12 valence f. Se 2 has 6 + 2(7) = 20 valence electrons. electrons. Se g. 2 has 4 + 2(6) = 16 valence electrons h. 2 has 2(6) = 12 valence electrons.

21 258 APTER 8 BDIG: GEERAL EPTS i. Br has = 8 valence electrons. Br 68. a. P 3 has (7) = 32 valence electrons. P Skeletal structure P Lewis structure This structure uses all 32 e while satisfying the octet rule for all atoms. This is a valid Lewis structure. S 4 2 has 6 + 4(6) + 2 = 32 valence electrons. S 2- ote: A negatively charged ion will have additional electrons to those that come from the valence shell of the atoms. Xe 4, 8 + 4(6) = 32 e P 4 3, 5 + 4(6) + 3 = 32 e Xe 4 has 7 + 4(6) + 1 = 32 valence electrons 3- P

22 APTER 8 BDIG: GEERAL EPTS ote: All of these species have the same number of atoms and the same number of valence electrons. They also have the same Lewis structure. b. 3 has 5 + 3(7) = 26 valence electrons. S 3 2, 6 + 3(6) + 2 = 26 e Skeletal structure Lewis structure S 2- P 3 3, 5 + 3(6) + 3 = 26 e P 3-3, 7 + 3(6) + 1 = 26 e ote: Species with the same number of atoms and valence electrons have similar Lewis structures. c. 2 has 7 + 2(6) + 1 = 20 valence - Skeletal structure Lewis structure S 2, 6 + 2(7) = 20 e P 2, 5 + 2(7) + 1 = 20 e S P - ote: Species with the same number of atoms and valence electrons have similar Lewis structures. d. Molecules ions that have the same number of valence electrons and the same number of atoms will have similar Lewis structures. 69. Be 2, 2 + 2(1) = 4 valence electrons B 3, 3 + 3(1) = 6 valence electrons

23 260 APTER 8 BDIG: GEERAL EPTS Be B 70. a. 2, 5 + 2(6) = 17 e 2 4, 2(5) + 4(6) = 34 e plus others plus other resonance structures b. B 3, 3 + 3(1) = 6 e 3, 5 + 3(1) = 8 e B B 3 3, = 14 e B In reaction a, 2 has an odd number of electrons so it is impossible to satisfy the octet rule. By dimerizing to form 2 4, the odd electron on two 2 molecules can pair up, giving a species whose Lewis structure can satisfy the octet rule. In general, odd electron species are very reactive. In reaction b, B 3 is electron deficient. Boron has only six electrons around it. By forming B 3 3, the boron atom satisfies the octet rule by accepting a lone pair of electrons from 3 to form a fourth bond. 71. P 5, 5 +5(7) = 40 valence electrons S 4, 6 + 4(7) = 34 e P S 3, 7 + 3(7) = 28 e Br 3, 3(7) + 1 = 22 e

24 APTER 8 BDIG: GEERAL EPTS 261 Br Br Br Row 3 and heavier nonmetals can have more than 8 electrons around them when they have to. Row 3 and heavier elements have empty d orbitals which are close in energy to valence s and p orbitals. These empty d orbitals can accept extra electrons. or example, P in P 5 has its five valence electrons in the 3s and 3p orbitals. These s and p orbitals have room for 3 more electrons, and if it has to, P can use the empty 3d orbitals for any electrons above S 6, 6 + 6(7) = 48 e 5, 7 + 5(7) = 42 e S Xe 4, 8 + 4(7) = 36 e Xe 73. a. 2 has 5 + 2(6) + 1 = 18 valence electrons. The skeletal structure is: To get an octet about the nitrogen and only use 18 e -, we must form a double bond to one of the oxygen atoms. Since there is no reason to have the double bond to a particular oxygen atom, we can draw two resonance structures. Each Lewis structure uses the correct number of electrons and satisfies the octet rule, so each is a valid Lewis structure. Resonance structures occur when you have multiple bonds that can be in various positions. We say the actual structure is an average of these two resonance structures. 3 has 5 + 3(6) + 1 = 24 valence electrons. We can draw three resonance structures for 3, with the double bond rotating among the three oxygen atoms.

25 262 APTER 8 BDIG: GEERAL EPTS 2 4 has 2(5) + 4(6) = 34 valence electrons. We can draw four resonance structures for 2 4. b. has = 16 valence electrons. We can draw three resonance structures for. S has = 16 valence electrons. Three resonance structures can be drawn. S S S 3 has 3(5) + 1 = 16 valence electrons. As with and S, three different resonance structures can be drawn. 74. zone: 3 has 3(6) = 18 valence electrons.

26 APTER 8 BDIG: GEERAL EPTS 263 Sulfur dioxide: S 2 has 6 + 2(6) = 18 valence electrons. S S Sulfur trioxide: S 3 has 6 + 3(6) = 24 valence electrons. S S S 75. Benzene has 6(4) + 6(1) = 30 valence electrons. Two resonance structures can be drawn for benzene. The actual structure of benzene is an average of these two resonance structures, that is, all carbon-carbon bonds are equivalent with a bond length and bond strength somewhere between a single and a double bond. 76. Borazine (B ) has 3(3) + 3(5) + 6(1) = 30 valence electrons. The possible resonance structures are similar to those of benzene in Exercise B B B B B B

27 264 APTER 8 BDIG: GEERAL EPTS 77. We will use a hexagon to represent the six-member carbon ring, and we will omit the 4 hydrogen atoms and the three lone pairs of electrons on each chlorine. If no resonance existed, we could draw 4 different molecules: If the double bonds in the benzene ring exhibit resonance, then we can draw only three different dichlorobenzenes. The circle in the hexagon represents the delocalization of the three double bonds in the benzene ring (see Exercise 8.75). With resonance, all carbon-carbon bonds are equivalent. We can t distinguish between a single and double bond between adjacent carbons that have a chlorine attached. That only 3 isomers are observed supports the concept of resonance has 4 + 3(6) + 2 = 24 valence electrons Three resonance structures can be drawn for 3 2. The actual structure for 3 2 is an average of these three resonance structures. That is, the three bond lengths are all equivalent, with a length somewhere between a single and a double bond. The actual bond length of 136 pm is consistent with this resonance view of (10 e -): Triple bond between and. 2 4 (38 e -) : Single bond between and. 2 2 (24 e -) : Double bond between and.

28 APTER 8 BDIG: GEERAL EPTS 265 As the number of bonds increase between two atoms, bond strength increases and bond length decreases. rom the Lewis structure, the shortest to longest - bonds are: 2 < 2 2 < The Lewis structures for the various species are: (10 e -): (10 e Triple bond between and ): Triple bond between and. 2 (16 e 2 (16 e -):): Double Double bond between bond between and and (24 e -): Average of 1 1/3 bond between and in (14 e ): Single bond between and. As the number of bonds increases between two atoms, bond length decreases and bond strength increases. With this in mind, then: ormal harge longest shortest bond: 3 > 3 2 > 2 > weakest strongest bond: 3 < 3 2 < 2 < 81. See Exercise 8.68a for the Lewis structures of P 3, S 4 2, 4 and P 4 3. ormal charge = [number of valence electrons on free atom] - [number of lone pair electrons on atom + 1/2 (number of shared electrons of atom)]. a. P 3 : P, = 5 1/2(8) = +1 b. S 4 2 : S, = 6 1/2(8) = +2 c. 4 :, = 7 1/2(8) = +3 d. P 4 3 : P, = 5 1/2(8) = +1 e. S 2 2, 6 + 2(6) + 2(7) = 32 e f. Xe 4, 8 + 4(6) = 32 e -

29 266 APTER 8 BDIG: GEERAL EPTS S Xe S, = 6 1/2(8) = +2 Xe, = 8 1/2(8) = +4 g. 3, 7 + 3(6) + 1 = 26 e h. 3 4, 5 + 4(6) + 3 = 32 e 3-, = 7 2 1/2(6) = +2, = 5 1/2(8) = or S 4 2, 4, P 4 3 and 3, only one of the possible resonance structures is drawn. a. Must have five bonds to P to minimize b. Must form six bonds to S to minimize formal charge of P. The best choice is formal charge of S. to form a double bond to since this will give a formal charge of zero and single bonds to for the same reason. P P, = 0 S 2- S, = 0 c. Must form seven bonds to d. Must form five bonds to P to to minimize formal charge. to minimize formal charge., = P 3- P, = 0 e. f. S, = 0 S, = 0, = 0 Xe Xe, = 0

30 APTER 8 BDIG: GEERAL EPTS 267 g. -, = 0 h. We can t. The following structure has a zero formal charge for : 3- But does not expand its octet. We wouldn t expect this resonance form to exist has 2(6) + 2(7) = 26 valence e. The formal charge and oxidation number of each atom is below the Lewis structure of 2 2. ormal harge xid. umber xidation numbers are more useful when accounting for the reactivity of 2 2. We are forced to assign +1 as the oxidation number for oxygen. xygen is very electronegative, and +1 is not a stable oxidation state for this element. 84. has = 16 valence electrons. ormal charge nly the first two resonance structures should be important. The third places a positive formal charge on the most electronegative atom in the ion and a 2 formal charge on. will also have 16 valence electrons. ormal charge

31 268 APTER 8 BDIG: GEERAL EPTS All of the resonance structures for fulminate ( ) involve greater formal charges than in cyanate ( ), making fulminate more reactive (less stable). 85. S, = 13; the formula could be S (13 valence electrons), S 2 2 (26 valence electrons), S 3 3 (39 valence electrons), etc. or a formal charge of zero on S, we will need each sulfur in the Lewis structure to have two bonds to it and two lone pairs [ = 6 4 1/2(4) = 0]. will need one bond and three lone pairs for a formal charge of zero [ = 7 6 1/2(2) = 0]. Since chlorine wants only one bond to it, it will not be a central atom here. With this in mind, only S 2 2 can have a Lewis structure with a formal charge of zero on all atoms. The structure is: S S 86. The nitrogen-nitrogen bond length of 112 pm is between a double (120 pm) and a triple (110 pm) bond. The nitrogen-oxygen bond length of 119 pm is between a single (147 pm) and a double bond (115 pm). The third resonance structure shown below doesn t appear to be as important as the other two since there is no evidence from bond lengths for a nitrogen-oxygen triple bond or a nitrogen-nitrogen single bond as in the third resonance form. We can adequately describe the structure of 2 using the resonance forms: Assigning formal charges for all 3 resonance forms: or: , = /2(4) = -1, = 5-1/2(8) = +1, Same for and, = /2(2) = -2 ;, = /2(6) = 0, = /2(4) = 0 ;, = /2(2) = -1, = /2(6) = +1

32 APTER 8 BDIG: GEERAL EPTS 269 We should eliminate since it has a formal charge of +1 on the most electronegative element (). This is consistent with the observation that the bond is between a double and triple bond and that the bond is between a single and double bond. Molecular Structure and Polarity 87. The first step always is to draw a valid Lewis structure when predicting molecular structure. When resonance is possible, only one of the possible resonance structures is necessary to predict the correct structure because all resonance structures give the same structure. The Lewis structures are in Exercises 8.67 and The structures and bond angles for each follow a. : linear, 180 b. P 3 : trigonal pyramid, < c. 3 : tetrahedral, d. 4 + : tetrahedral, e. 2 : trigonal planar, 120 f. Se 2 : V-shaped or bent, < g. 2 : linear, 180 h and i. 2 and Br are both linear, but there is no bond angle in either. ote: P 3 and Se 2 both have lone pairs of electrons on the central atom which result in bond angles that are something less than predicted from a tetrahedral arrangement (109.5 ). owever, we cannot predict the exact number. or the solutions manual, we will insert a less than sign to indicate this phenomenon. or bond angles equal to 120, the lone pair phenomenon isn t as significant as compared to smaller bond angles. or these molecules, e.g., 2, we will insert an approximate sign in front of the 120 to note that there may be a slight distortion from the VSEPR predicted bond angle a. 2 : V-shaped, 120 ; 3 : trigonal planar, : trigonal planar, 120 about both atoms b., S and 3 are all linear with 180 bond angles. 88. See Exercises 8.68 and 8.74 for the Lewis structures a. All are tetrahedral; b. All are trigonal pyramid; < c. All are V-shaped; < and S 2 are V-shaped (or bent) with a bond angle 120. S 3 is trigonal planar with 120 bond angles.

33 270 APTER 8 BDIG: GEERAL EPTS 89. rom the Lewis structures (see Exercise 8.71), Br 3 would have a linear molecular structure, 3 would have a T-shaped molecular structure and S 4 would have a see-saw molecular structure. or example, consider 3 (28 valence electrons): The central atom is surrounded by 5 electron pairs, which requires a trigonal bipyramid geometry. Since there are 3 bonded atoms and 2 lone pairs of electrons about, we describe the molecular structure of 3 as T-shaped with predicted bond angles of about 90. The actual bond angles would be slightly less than 90 due to the stronger repulsive effect of the lone pair electrons as compared to the bonding electrons. 90. rom the Lewis structures (see Exercise 8.72), Xe 4 would have a square planar molecular structure, and 5 would have a square pyramid molecular structure. 91. a. Se 3, 6 + 3(6) = 24 e o Se 120 o Se Se 120 o Se 3 has a trigonal planar molecular structure with all bond angles equal to 120. ote that any one of the resonance structures could be used to predict molecular structure and bond angles. b. Se 2, 6 + 2(6) = 18 e - Se 120 o Se 2 has a V-shaped molecular structure. We would expect the bond angle to be approximately 120 as expected for trigonal planar geometry. ote: Both of these structures have three effective pairs of electrons about the central atom. All of the structures P are based on a trigonal planar geometry, but only Se 3 is described as having a trigonal planar structure. Molecular structure always describes the relative positions of the atoms. 92. a. P 3 has 5 + 3(7) = b. S 2 has 6 + 2(7) = 26 valence electrons. 20 valence electrons. Se S Trigonal pyramid; all angles are < V-shaped; angle is <

34 APTER 8 BDIG: GEERAL EPTS 271 c. Si 4 has 4 + 4(7) = 32 valence electrons. Tetrahedral; all angles are ote: There are 4 pairs of electrons about the central atom in each case in this exercise. All of the structures are based on a tetrahedral geometry, but only Si 4 has a tetrahedral structure. We consider only the relative positions of the atoms when describing the molecular structure. 93. a. Xe 2 has 8 + 2(7) = 22 valence electrons. Si There are 5 pairs of electrons about the central Xe atom. The structure will be based on a trigonal bipyramid geometry. The most stable arrangement of the atoms in Xe 2 is a linear molecular structure with a 180 bond angle. b. I 3 has 7 + 3(7) = 28 valence electrons. Xe 180 o I 90 o 90 o T-shaped; The I angles are 90. Since the lone pairs will take up more space, the I bond angles will probably be slightly less than 90. c. Te 4 has 6 + 4(7) = 34 d. P 5 has 5 + 5(7) = 40 valence electrons. valence electrons. 120 o Te 90 o See-saw or teeter-totter or distorted tetrahedron 120 o Trigonal bipyramid All of the species in this exercise have 5 pairs of electrons around the central atom. All of the structures are based on a trigonal bipyramid geometry, but only in P 5 are all of the pairs bonding pairs. Thus, P 5 is the only one we describe as a trigonal bipyramid molecular structure. Still, we had to begin with the trigonal bipyramid geometry to get to the structures of the others. P 90 o 94. a. I 5, 7 + 5(7) = 42 e - b. Xe 4, 8 + 4(7) = 36 e - 90 o 90 o I o Xe 90 o 90 o