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1 hem 1A practice problems 4/5/16 (x-track) 1. Write the condensed electron configuration for each element from Be to. ow many valence electrons do they have? Repeat with each element from Si to l. Be: [e]2s 2, : [e]2s 2 2p 5 Si: [e]3s 2 3p 2, l: [e]3s 2 3p 5 Be: 2, B: 3, : 4, : 5, : 6, : 7 Si: 4, P: 5, S: 6, l 7 2. ow many electrons fit in each orbital? Why? What must be true of electrons in the same orbital? 2 electrons per orbital, and they must have opposite spins. Pauli exclusion principle says that no two electrons in the same atom can have all the same quantum numbers, so if they are in the same orbital, 3 Q are the same, so the last one must be different. S S S I P P P Kr Kr Si Se l 3 S 2+ l S
2 3. Draw Lewis dot structures of each of the following (use lines for bonds and dot for nonbonding electrons): a) 2 b) S c) 26 d) 3 e) S6 f) S2 g) I3 h) P4 3 See image above. 4. Draw Lewis structures: B3, S8, S2, Pl3 5. Use Lewis structures to decide which has the longest and shortest - bond: 2, 22, 24. See first image. As we add, the number of bonds between the s goes down, so the bond length should increase. Triple bond, shortest; single bond, longest. 6. What is the difference between formal charges and oxidation numbers? When do you use each one? ormal charges are a bookkeeping method used when drawing Lewis structures. When counting formal charges, we assume bonding electrons are divided evenly between the two atoms. xidation numbers are a bookkeeping method used when deciding if a redox reaction has occurred. When counting oxidation numbers, we assume that bonding atoms belong only to the atom that attracts them more. 7. Draw Lewis structures, including formal charges, for each. a) 4 + b) c) l4 d) S3 See first image. 8. Draw Lewis structures, showing bond dipole moments. a) 32 b) l c) 2S See first image. 9. Predict the electron domain geometry and molecular geometry of each molecule you drew a Lewis structure for. a) 2: no shape (only 2 atoms) b) S : linear c) 26: each is tetrahedral d) 3: trigonal pyramidal e) S6: octahedral f) S2: bent g) I3: T-shape h) P4 3 : tetrahedral carbonate and nitrate are trigonal planar. 22 is bent at each. 24 is trigonal pyramidal at each.
3 10., 2, and 3 are polar molecules. 4 and B3 are not polar. What are the shapes of these molecules? ow do you know? Does this match the prediction based on the Lewis structures and VSEPR? Water and ammonia are polar because they have lone pairs, and thus the bond dipoles don t cancel out. Water is bent, and ammonia is trigonal pyramidal. Methane is tetrahedral, and thus not polar. Borane is trigonal planar, and not polar. either has lone pairs. You can predict the shapes using the electron domain model, and confirm using the polarity data. 11. Draw Lewis structures (with formal charges) and predict the shape of the molecule. a) Xe2 b) S4 c) I3 geometry: trigonal bipyramidal. 1 lone pair: seesaw c) trigonal bipyramidal, 2 lone pairs, T-shape a)electron domain geometry: trigonal bipyramidal. 3 lone pairs take equatorial positions, so it is linear. b)electron domain 12. Make a table: for each element from Li to e, how many bonds does it make, and how many lone pairs does it have, when it has zero formal charge? (int: is there a limit on how many electrons these elements can have total? What is it?) Memorize this table and use it to help you draw good Lewis structures quickly. Be: 2 bonds, no lone pairs B: 3 bonds, no lone pairs : 4 bonds, no lone pairs : 3 bonds, 1 LP : 2 bonds, 2 LP : 1 bond, 3 LP e: 0 bonds, 4 LP 13.Draw Lewis structures (including resonance structures) for 3, 2, 3, and S23 2. otice that in each case, some bonds are multiple bonds and other bonds are single bonds, but there s nothing different about the atoms making the single and double bonds, no reason why
4 1 they should be different. Experimentally, we find that the bonds aren t different, which means we have to draw multiple resonance structures to show that the bonds are the same. 14. Sometimes you might forget what the charge on carbonate, nitrate, or other ions are. Use Lewis structures to decide what charges should be stable for 3 n and 3 n. See first image. They really tell you odd/even charge. carbonate could be 2- or 4-; nitrate could be 1- or 3-. In both cases, the lower charge is correct. 15. What is the meaning of resonance structures? Resonance structures are a way to deal with situations that Lewis structures can t handle well, like 3. In reality, there are strength bonds, or 2 single bonds a a double bond shared over 3 atoms. But we can t really show this well with Lewis structures except by drawing the double bond in both places using resonance structures. 16. or each pair of molecules, which do you think would be hardest to make, based on Lewis structures? a) P3 vs P4 b) 2 vs Kr2 c) Si2 vs Se2 Refer to the image. a) otice that P4 has a radical (an unpaired electron). This will tend to make the molecule unstable. P3 is fine. b) We can t tell yet which of the 2 structures is most appropriate, but at least it looks like 2 is more stable than atoms in the gas phase. Probably some bonds will form. Kr has no reason to form bonds, because it already has octet, so it would be very hard to get 2 Kr atoms to stick to each other. c) Se2 is a fine Lewis structure. Si2 is a problem because Si doesn t have octet, so it would probably be hard to make and very reactive. 17. What is the difference between l and [ + ][l ]? Which is a better description of the molecule, or is neither exactly right? ne is the covalent Lewis structure, indicating that a pair of electrons is shared between l and. The other is the ionic Lewis structure, indicating a chloride ion and a hydrogen ion, connected by ionic attraction. The best (most complete) description is probably both, indicating that the bond is part ionic and part covalent. (y-track) 18. Mulliken defined electronegativity as proportional to the sum of electron affinity and ionization energy. Why does this make sense? Refer to the data in the first table. Yes, this does make sense. Electronegativity is the ability of an atom to pull electrons toward itself, and resist other atoms pulling its electrons away. If EA is big, it can pull electrons toward itself. If IE is big, it can resist other atoms pulling its electrons away. Looking at the data, we should find that E (electronegativity) is biggest in the top right of the periodic table, which makes sense.
5 19. Look at the data in the second table for Dxx bond enthalpies. What patterns in bond strength do you notice? Where in the periodic table are bonds stronger and weaker? -, -, and - single bonds are weak. - and - single bonds are strong. Multiple bonds are strong. 20. Look at the data in the third table, which shows bond enthalpies for X-Y bonds and compares them to the average of the X-X and Y-Y bond strengths. Does the resonance energy depend on the polarity of the bond? Yes, it seems like more polar bonds have bigger resonance energy. (Remember bonds are more polar if the electronegativity (E) is different between the two atoms. E is biggest in the top right of the periodic table ( and ). Moving down or left, E decreases.) 21. Resonance (like in Q20, and in Q13) makes molecules more stable when there are more good (or ok) Lewis structures. (The more possible structures, and the more reasonable each is, the lower the total energy of the molecule.) ow would you write the resonance involved in Q20 as Lewis structures? Why does the resonance energy in Q20 vary as it does? See Q17. You could write ionic resonance structures. The bonds get stronger as the ionic resonance structure gets more reasonable. bond l-l = 144 I-I = Pauling defined electronegativity using the resonance energy, as calculated in the third table. The difference in E between X and Y is proportional to the resonance energy. Does this make sense? Explain. Yes, see previous enthalpy (kcal/mol) bond Dxy (Dxx + Dyy)/2 resonance energy Dxy (Dxx + Dyy)/ I l l l l-i I I element IE (ev) EA (ev) element IE (ev) EA (ev)
6 l ± 0.2 Br I a B u Au
2. Atoms with very similar electronegativity values are expected to form
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