CHEMISTRY 333 Fall 2006 Exam 3, in class

Transcription

1 CHEMISTRY 333 Fall 006 Exam 3, in class Name SHOW YOUR WORK. Mathematica & calculators. You may use Mathematica or a calculator to do arithmetic. Do not use Mathematica for any higher-level manipulations (algebra, equation solving and/or simplification, integration, differentiation, eigensystem calculations, etc.). SHMO, Textbook, Lecture notes, etc.. No SHMO usage is allowed. The exam is closed-book, closedweb page, closed lecture notes, closed homework assignments, etc. 1a. What values of the n, l, and m l quantum numbers characterize the 3d z wave function? (Hint: the 3d z wave function is a function of r and θ only; it is not a function of φ). n = 3, l = (d orbital), m l = 0 (not a function of φ) 1b. How many hydrogen atom wave functions are degenerate with ψ 310? List them. Hydrogen atom orbital energies depend on n only. There are 9 (n ) orbitals with n = 3 so 8 orbitals are degenerate with 310: 300, 31-1, 311, 3-, 3-1, 30, 31, 3. 1c. What boundary condition(s) are applied to the radial component of the hydrogen atom wave function? What consequences do these conditions have for the physical properties of a hydrogen atom? Boundary condition wave function approaches zero as distance (r) becomes large. This condition, like all other boundary conditions that we have seen, results in the energy being quantized.

2 Recall that [ Ĥ, ˆl ] = [ l ˆz, ˆl ] = 0 for the hydrogen atom. Also, ˆ ml l = ( + 1) ly ll Y and ˆ ml ml ly = my. ml l z l l l. Evaluate the following integrals. Explain your reasoning (no integrations are necessary). ψ Ĥψ dτ This integral equals zero. The 10 and 310 functions are eigenfunctions of the Hamiltonian and are orthogonal. Going a bit farther the integrand can be rewritten as Ψ 310 E 10 Ψ 10, and the energy can be removed from the integrand because it is a constant, leaving the integrand as the product of orthogonal functions: Ψ 310 Ψ 10 ψ l ψ dτ 10 ˆz 10 This integral equals zero. Applying the operator to the 10 function multiplies the function by m l, but m l = 0 so the integrand vanishes. ( Y ) lˆ Y dτ 1 * 1 This integral equals zero. The spherical harmonics are eigenfunctions of the operator and are orthogonal. The reasoning is similar to that used in the first integral. 3. Three quantum chemists approach you for research funds for the same research idea: use the variation method to develop a model wave function of the hydrogen atom ground state. Chemist A wants to optimize φ(r) = N exp(-ar ). Chemist B wants to optimize φ(r) = N (ar + br + cr 6 ). Chemist C wants to optimize φ(r) = N sin(ar). The variational parameters in these functions are a, b, and c. You only have money to fund one project. Which one do you fund and why? (Hint: Explain why two of the ideas are definitely worthless.) Fund chemist A. The others are wasting their time. The variation method provides a way for optimizing model wave functions only when the model observes the same mathematical and physical constraints observed by the real wave function. Chemist B wants to optimize a model that blows up when r gets large (the real wave function vanishes). Similarly, Chemist C wants to optimize a model that fails to decay as r gets large. Only Chemist A has a model that obeys the constraints on the system. Note: don t mistake A s model for the real wave function which is N exp(-ar).

3 3 a. The first ionization potentials (IP) of H, He, and Li are: 13.6 (H),.6 (He), and 5. (Li) ev. The nuclear charge (Z) in these atoms, however, increases +1 (H), + (He), and +3 (Li). Explain the up-and-down trend in IP in terms of specific interactions and quantum mechanical principles. In the independent particle approximation, the IP is the negative of the electron s orbital energy. The energy of an orbital in a one-electron atom with nuclear charge Z depends on only two factors: n and Z. The energy increases with increasing Z and decreases with increasing n. The ionized electrons in H and He both come from 1s orbitals (n = 1). The dominant factor, then, is Z which is larger for He. As expected, He has a larger IP. The ionized electron in Li comes from a s orbital (n = ), which argues for smaller IP smaller. On the other hand, Z = 3, so this argues for larger IP. How to decide? The core (1s) electrons shield the s electron from the nucleus so, in fact, Z eff is between 1 and, and the change in n dominates. In terms of specific interactions, shielding is created by electron-electron repulsions that offset nuclear-electron attraction. b. The nuclear charge on Be (Z = +) is larger than that of Li (Z = +3). Describe whether Be s first IP is larger or smaller than Li s. Justify your answer in terms of specific interactions and quantum mechanical principles. Li has a 1s s 1 electron configuration and Be has a 1s s electron configuration. n = for the ionized electron in both cases. Z and also Z eff are larger for Be than for Li, so Be s IP should be larger. Remember that shielding, electron-electron repulsion that offsets nuclear-electron attraction, is mainly created by inner shell electrons. Li and Be both have two inner electrons. Electrons in the same orbital repel, but this repulsion does not offset nuclear-attraction that much.

4 5. Write the Hamiltonian operator in for H assuming fixed nuclei, i.e., do not include operators for nuclear kinetic energy. To make the operators easier to write, use Cartesian coordinates for the kinetic energies, but use interparticle distances for potential energies. There are four particles in H. Thus, there should be kinetic energy operators and 6 potential energy operators (each particle-particle interaction is described by a separate operator). The kinetic energy operators for the nuclei can be omitted because we assume fixed nuclei. Particle coordinates are indicated as follows: 1 = electron 1, = electron, A = nucleus A, and B = nucleus B. where ˆ H = ( 1 + ) me πε o r1a r1b ra rb r1 rab i = + + xi yi zi A chemist decides to use the following two-determinant model wave function for H. (Ignore normalization.) A refers to an atomic orbital on one H and B refers to an atomic orbital on the other H. A(1) α(1) B(1) β(1) B(1) α(1) A(1) β(1) Ψ (1, ) = + A() α() B() β() B() α() A() β() 6a. Rewrite Ψ(1,) as the product of a two-electron spatial function and a two-electron spin function. No one had trouble expanding these determinants, but several folks stopped after expanding the function. They did not factor apart the spatial and spin functions. I wanted the following, a product of two functions, one function that depends on spatial coordinates only and another that depends on spin coordinates only: [ A(1) B() B(1) A() ][ α(1) β() β(1) α() ] Ψ= +

5 5 6b. What are the symmetries of the spatial and spin functions in part 6a with respect to electron interchange? Does Ψ(1,), the combined function, satisfy the Pauli principle? Explain. Interchanging the coordinates of electrons 1 and gives: [ A() B(1) + B() A(1) ][ α() β(1) β() α(1) ] The spatial function is unchanged ( symmetric ) from 6a. The spin function is the negative of 6a ( antisymmetric ). Overall, the wave function is the negative of the original ( antisymmetric ), so it satisfies the Pauli principle which requires a wave function to change sign when the coordinates of any two electrons are interchanged. 6c. Is Ψ(1,) an eigenfunction of S ˆz? Show your work. If it is an eigenfunction, what is M S? Possibly helpful formulas: Sˆ = sˆ + sˆ sˆ α = ( /) α sˆ β = ( /) β z 1z z z In working this problem, it is helpful to notice that a spin operator works on spin functions only, so the spatial component of Ψ can be ignored. As a further simplification, I will use k to represent /. The following algebra shows that the operator makes the wave function vanish. In other words, the wave function is an eigenfunction with an eigenvalue of zero (M S = 0). S` z HaH1L bhl - bh1l ahll = = Hs`1 z + s` zl HaH1L bhl - bh1l ahll = s`1 z HaH1L bhl - bh1l ahll + s` zhah1l bhl - bh1l ahll = k a(1)b()-(-k)b(1)a() + (-k)a(1)b()-(k)b(1)a() = 0 z

6 6 6d. Is Ψ(1,) an eigenfunction of Ŝ? Show your work. Eigenfunctions of ˆ S ψ S = S( S + 1) ψ S are characterized by a quantum number S. If Ψ(1,) is an eigenfunction of is its S quantum number? Possibly helpful formulas (see part 6b too): Ŝ, like ψ S in the formula Ŝ, what Sˆ = sˆ + sˆ + ( sˆ sˆ + sˆ sˆ + sˆ sˆ ) 1 1x x 1y y 1z z sˆ α = (3 / ) α sˆ β = (3 / ) β sˆ α = ( /) β sˆ β = ( /) α sˆ α = ( i /) β sˆ β = ( i /) α x x This is only a little more complicated than the previous problem. Once again we can (and should) ignore the spatial function. I ll leave some of the math details to you and just point out the following: y y s ` 1 HaH1L bhl -bh1l ahll = s ` HaH1L bhl -bh1l ahll = 3 hbar 3 hbar HaH1L bhl -bh1l ahll HaH1L bhl -bh1l ahll s`1 x s` x HaH1L bhl -bh1l ahll = s`1 y s` y HaH1L bhl -bh1l ahll = hbar hbar H bh1l ahl -ah1l bhll H bh1l ahl -ah1l bhll s`1 z s` z HaH1L bhl -bh1l ahll = - hbar HaH1L bhl - bh1l ahll = hbar H bh1l ahl -ah1l bhll Combining all of these terms shows that the operator makes the wave function vanish. In other words, the wave function is an eigenfunctions of the operator with an eigenvalue of zero (S = 0).

7 7 The following graphs are based on a particular hydrogen atomic orbital Graph # Graph # a. Which graph shows ψ and which shows the radial distribution function of ψ. Explain your reasoning. I believe that I was asked during the exam whether I really meant Ψ and I said to change this to Ψ. Some of you used this information to answer the question, but others did not. Regardless, the radial distribution function is the one on the right. The radial distribution function is proportional to r Ψ, so it rises more slowly near the nucleus (where r is small) and its maximum occurs at larger r. 7b. For each atomic orbital listed below, explain why the listed orbital might or might not be the orbital responsible for the graphs. (Note: Some explanations may be re-usable.) 1s The graphs cannot refer to this orbital because a 1s orbital has its maximum value at r = 0 (Note: the 1s orbital is not infinite anywhere.) s The graphs cannot refer to this orbital because a s orbital has a non-zero value at r = 0. This is true of all s orbitals. p z The graphs might refer to this orbital. Its radial function vanishes at the nucleus only. 3p z The graphs cannot refer to this orbital because the 3p z orbital contains a radial node at r > 0. 3d xy The graphs might refer to this orbital. Its radial function vanishes at the nucleus only.

8 8 A chemist decides to model the wave function for H + at its equilibrium geometry as Ψ = N(ψ 1sA + ψ 1sB) where ψ 1si is the 1s atomic orbital on atom i. Assume the two nuclei are located on the z axis at -r/ and +r/, respectively. 8a. Sketch the value of Ψ as a function of z for H + for x = 0, y = 0, -r z +r A 1s orbital is a decaying exponential that takes on its maximum value at the nucleus. A linear combination of these functions looks something like the following where the two maxima are located approximately at r/ and +r/, respectively b. Another chemist suggests a more complicated model for H + that combines 1s and p z functions. Write an algebraic formula for Ψ complicated and identify all variational parameters. The labeling rule in 8a is to use the subscript to identify the type of atomic orbital and the atom that it is centered on. The complicated model combines atomic orbitals, from each atom: Ψ complicated = c 1 ψ 1sA + c ψ pza + c 3 ψ 1sB + c ψ pzb The complicated model wave function is formed by multiplying each atomic orbital function by a weighting coefficient, a number, that tells us how much the orbital contributes to the model. The four weighting coefficients are variational parameters because we adjust them to make the model s energy as low as possible. 8c. How would you use E original and E complicated, the energies of the two model wave functions, to decide which model is closer to the *** Quantum Mechanical Truth ***? Since the complicated model encompasses all of the features of the original model, it must do at least as well. This does not answer the question, however, as to how energies indicate model quality. The variation principle guarantees that model energies are always greater than the true energy, E model > E QMT. The closer E model comes to E QMT, the closer the model comes to QM truth. If the complicated model is really an improvement on the original model, we should observe E original > E complicated > E QMT.

9 9 9. The table shown above was discussed in class in connection with bonding in H +. Explain how these data were generated and how they shed light on the mechanism of covalent bonding in this ion. The energy values in the Total column were obtained using a model wave function that combines two 1s orbitals in a bonding fashion ( original wave function in problem 8). The energy of this model was calculated at two geometries well-separated nuclei and nuclei at the equilibrium bond distance and the energy difference between these geometries is shown in the Total column. In other words, when the two nuclei are pushed together, the kinetic energy rises, the potential energy falls by twice as much (satisfying the virial theorem), and the total energy falls too. Each value in the Total column is obtained using a particular energy operator (or operators) to calculate the energy. ΔT uses the electron kinetic energy operator. ΔV combines three interaction operators: two electron-nucleus interactions (the electron interacts independently with each nucleus) and one nucleusnucleus interaction. ΔE = ΔT + ΔV. (ΔE can also be obtained using the total energy operator, the Hamiltonian, for H +.) As discussed in class, the 1s orbitals use the same exponent, but this exponent is optimized for each geometry. The exponent s value when the nuclei are well-separated is identical to its value in the hydrogen atom. The exponent s value increases when the nuclei are separated by the bond distance. This compresses the orbital around each nucleus, and this affects the kinetic and potential energies. Since changes in orbital exponent and geometry both affect energy, a nonbonding model was constructed to separate the two effects. The nonbonding model, defined as ψ 1sA + iψ 1sB, combined the 1s orbitals in a fashion that delocalizes the electron, but doesn t produce bonding. Important: the model uses the same orbital exponents used above. The energy changes for this model are shown in the Nonbonding column. The kinetic energy rises strongly and the potential energy falls strongly. These effects are due mainly to orbital compression because the model is nonbonding (notice that the total energy change is small). The energy values in the Bonding column were obtained by subtracting the values in the Nonbonding column from the values in the Total column. The idea behind these data is that they show the leftover energy changes that occur when the nuclei approach, i.e., the energy changes separate from the changes caused by orbital compression. Supposedly, these leftovers are the changes responsible for bonding.

10 10 A five carbon cation and the properties of its two occupied Huckel MO are shown below (the orbital coefficients have been normalized). MO α +.1 β α β energy c c c c c a. Write the Huckel MO secular matrix for this cation (use my numbering scheme). The secular matrix means the matrix of coefficients generated by the secular equations. This matrix is shown below. Many of you, however, wrote one of the matrices that can be obtained by transforming the secular matrix and I allowed varying amounts of credit for these depending on their similarity to the secular matrix. i j k α E β β α E β 0 β 0 β α E β β α E β 0 β 0 β α E y z { 10b. Why do we insist that the determinant of the secular matrix equals zero? Because this provides the only physically meaningful solution to the secular equations. The equations all have a similar form, such as 0 = c 1 (α E) + c (β) + c c 0 + c 5 0 (this is first secular equation based on the matrix in 10a). One solution to these equations is to set c 1 = c = = c 5 = 0, but this is not physically meaningful. The only other solution is to require that the determinant of the secular matrix equal zero. 10c. Calculate E π for this cation. In Huckel theory, E π is the sum of the electron energies. The ion contains four electrons and these occupy the two lowest energy MOs (E 1 = α +.1 β, E = α β). E π = E 1 + E = α β

11 11 10d. Calculate the number of pi electrons on each carbon (see previous page for coefficients and numbering). Based on these results, assign a charge to each carbon. The spatial distribution of each electron is determined by the square of its MO. In Huckel theory, we ignore terms containing p i p j where i and j are different atoms. In other words, we say ψ i c 1 p 1 + c p + c 3 p 3 + c p + c 5 p 5. The electron density distribution for this electron puts c i electrons on atom i. Two electrons are described by ψ 1 and two electrons are described by ψ, where the MOs are: ψ 1 =.6 p p +.6 p 3 +. p +.6 p 5 ψ =.66 p 1 +. p -.18 p p -.18 p 5 The number of pi electrons on carbon 1 is (.6 ) + (.66 ) = 1. Carbons and also hold 1 pi electron each. The charge on each of these three carbons is zero. The number of pi electrons on carbon 3 (and also carbon ) is (.6 ) + (-.18 ) = 0.5. The charge on each of these two carbons is e. Calculate the pi bond orders between each pair of neighboring carbons (skip bonds that are symmetry-related to other bonds) Huckel theory says the bond order created by one electron between atoms i and j is equal to c i c j where these are the coefficients from the electron s MO. To get the total bond order between atoms i and j we must add up the contributions from each electron. Bond order 1- = (.6 x.56) + (.66 x.) = 0.86 Bond order -3 = (.56 x.6) + (. x (-.18)) = 0.36 = Bond order -5 (by symmetry) Bond order 3- = (. x.6) + ((-.56) x (-.18)) = 0.61 = Bond order -5 (by symmetry)

12 1 The relative pi electron energies of anthracene and phenanthrene can be compared using Dewar PMO by uniting two benzyl radicals. vs. anthracene phenanthrene 11a. Calculate numerical values for the orbital coefficients in benzyl radical s NBMO using Dewar s rules. (You may not use SHMO for any part of problems 11 and 1; you may use Mathematica to perform arithmetic operations like multiplication-addition-square root, but nothing more complicated). You probably should have obtained the coefficients in three stages. First stage: non-zero coefficients are located on starred atoms only and the sums of the coefficients around any blank atom must equal zero. Second stage: in a normalized NBMO, the sums of the squares of the coefficients equal one. Third stage: numerical values for the coefficients. -1 (-1/7) 1/ (/7) 1/ (1/7) 1/ (-1/7) 1/ b. Calculate the change in NBMO energy that occurs when two benzyl radicals unite to form anthracene by using Dewar s PMO method. This could have been worded better. The NBMO energy does not change. Rather, combining the radicals produces new MO with different energies from the NBMO. What I want, of course, is the difference between the NBMO and new MO energies. As shown in the drawing at the top of the page, uniting two radicals to make anthracene brings together a benzylic carbon in one radical with an ortho carbon in another radical. The product of these coefficients is -.38 x.76 = -.9. Since this happens in two locations simultaneously, the new MO energy is lower than the NBMO energy by x.9 β 0.6 β. The overall change in pi energy is twice this because two electrons occupy the new MO. Some of you were confused by the opposite signs of the coefficients and predicted that the NBMO interaction would be destabilizing. In fact, it is both stabilizing and destabilizing. One new MO is a bonding combination of the NBMO; the other is an antibonding combination. I am interested in the bonding combination because this is the MO occupied by the two electrons. 11c. Repeat part 11b, but for benzyl phenanthrene. Phenanthrene is obtained by bringing the benzyl carbons together (.76 x.76) and the ortho carbons together (-.38 x -.38), so the change in NBMO energy is [.76 x.76 + (-.38 x -.38)] β =.7 β. The overall change in pi energy is twice this because two electrons occupy the new MO. 11d. Which molecule should have a more stable pi system, anthracene or phenanthrene, and by how much? Phenanthrene is expected to have a more stable pi system by (.7 -.6) β =.9 β (if you use lots of sig figs)

13 13 1XC. Find another way to use Dewar s PMO method to compare the energies of anthracene and phenanthrene and carry out the necessary calculations (show your work for each step: draw your fragments, calculate their NBMO coefficients, calculate the change in NBMO energy as the fragments unite, determine the relative pi energies of anthracene and phenanthrene). Do you obtain the same relative energies? Just in case I am not being clear, I want you to draw two fragments A and B that will make anthracene when united one way, and will make phenanthrene when united in another. A and B are identical benzyl radicals in problem 11, but it is not necessary (or even convenient) for A and B to be identical in this problem. This can be done many ways. You just have to make sure that your fragments are odd-alternant hydrocarbons and that the same fragments can give both anthracene and phenanthrene. Strange but true fact: some fragment choices predict that anthracene is more stable than phenanthrene. (This is very annoying, but come see me before you decide Dewar PMO is worthless b/c there may be a simple explanation for this paradox.) One way C vs. C Another way vs. And still another way (this one is cute because it uses two identical radicals, just like benzyl case) vs.