Magnetic Behavior and Electron Configuration of Compounds Version: 6.3

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1 Magnetic Behavior and Electron Configuration of Compounds Version: 6.3 Michael J. Vitarelli Jr. Department of Chemistry and Chemical Biology Rutgers University, 610 Taylor Road, Piscataway, NJ I. INTRODUCTION There are many forms of magnetism including: paramagnetism, diamagnetism, and ferromagnetism. More accurately, these are effects experienced by materials under the influence of a magnetic field. Ferromagnetism is the form that individuals are most familiar with. A material, such as iron, is strongly attracted to a magnetic field. This strong attraction is known as ferromagnetism. Similar to ferromagnetism, paramagnetism is an attraction to an external magnetic field; however, not nearly as strong as ferromagnetism. Paramagnetism occurs when a material has unpaired electrons. As a divergence from these, diamagnetism is a weak repulsion to a magnetic field. Diamagnetism occurs when all the electrons are paired in a material. No matter what the electron configurations are, all compounds are influenced at least to some extent by a magnetic field. In determining the electron configuration of a compound there are a number of rules that must be followed. First, electrons always occupy the lowest energy orbital available. In general, an electron in an orbital in a lower principal shell is at a lower energy state than an electron in an orbital in a higher principal shell. Furthermore, an electron in an s orbital is at a lower energy state than an electron in a p orbital, then an electron in a d orbital, than an electron in an f orbital; assuming all these orbitals are in the same principal shell. This tells us the relative energies of orbitals, but not how to fill them. Each principal shell contains a number of subshells, which contain a number of orbitals, and each orbital can hold a maximum of two electrons. Let s consider two examples of this. The first principal shell contains only s orbitals, thus only an s subshell. Each s subshell has one s orbital which can hold two electrons. Therefore, the first principal shell can hold only two electrons. The second principal shell can hold s orbitals and p orbitals, thus contains s subshells and p subshells. While s subshells can contain only one s oribtal, p subshells contain three p orbitals, again each orbital with two electrons. Thus p subshells can hold 6 electrons. Therefore the second principal shell can hold 8 electrons. This can be extended indefinitely; refer to Table I for a summary of principal shells and subshells. TABLE I: Principal Shells and Subshells Principal Quantum Number : n Angular Momentum Quantum Number : l 0 0,1 0,1,2 0,1,2,3... Subshells s s,p s,p,d s,p,d,f... Number of Orbitals in specific Subshell 1 1,3 1,3,5 1,3,5,7... Subshell Electron Capacity 2 2,6 2,6,10 2,6,10,14... Principal Shell Electron Capacity We ve learned that electrons occupy orbitals in increasing principal shell and increasing subshell designation, from s to p to d to f. However, we must now consider how we fill the subshells, specifically the orbitals in the subshells. We are filling these subshells with electrons, which are all negatively charged. We also know that like charges repel one another. Since electrons are negatively charged

2 they will repel each other. Thus, due to this electron-electron repulsion, it is more energetically favorable to fill empty orbitals first, then pair in occupied orbitals. Lastly, electrons can not occupy the same quantum state simultaneously. In other words, no two electrons can have the same four quantum numbers. This rule is known as the Pauli Exclusion Principle. The consequence being that if two electrons are in an orbital one electron is spin up and the other is spin down. As an example consider the electron configuration of carbon. Neutral carbon has 6 electrons. First, we add two electrons in the s orbital in the first principal shell. Then we add two electrons in the s orbital in the second principal shell. Finally, we add the last two electrons to the p subshell, still in the second principal shell. Recalling that p subshells have three p orbitals, these last two electrons will go into separate orbitals in this p subshell: Carbon : 1s 2s 2p 2p 2p Carbon : 1s 2 2s 2 2p 2 Notice carbon has two unpaired electrons. Thus carbon is a paramagnetic species; at least gas phase atomic carbon is paramagnetic. It s interesting to note that graphite and diamond, two allotropes of carbon, are both diamagnetic. Now let s consider a diamagnetic species. Neutral argon has 18 electrons. These electrons are placed in orbitals exactly how we would expect: Argon : 1s 2s 2p 2p 2p 3s 3p 3p 3p Argon : 1s 2 2s 2 2p 6 3s 2 3p 6 Again, since all the electrons in argon are paired the species is diamagnetic. Furthermore, notice the filling of the orbitals of argon follows all the rules we ve discussed above. Unfortunately, as the number of electrons in the species increases the filling of the atom s orbitals may become more complicated. Electrons always fill the lowest energy orbitals first, and we would expect that after the 3p orbitals we should fill the 3d orbitals. However, this is not the case. The 4s orbital is actually at a lower energy state than the 3d orbitals. Thus we fill the 4s orbital first, then the 3d orbitals. As an example consider the electron configuration of manganese: 2 Mn : 1s 2s 2p 2p 2p 3s 3p 3p 3p 4s 3d 3d 3d 3d 3d Mn : 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 Since manganese has 5 unpaired electrons manganese is paramagnetic. As our last example let s consider Zinc. Neutral Zinc has 30 electrons and has the following electron configuration: Zn : 1s 2s 2p 2p 2p 3s 3p 3p 3p 4s 3d 3d 3d 3d 3d

3 3 Zn : 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 Since zinc has all of it s electrons paired, zinc is diamagnetic. However, what happens if we try to ionize zinc? One might think that we would first remove the an electron from a 3d orbital, however this turns out to be not the case. When ionizing atoms we remove electrons from the orbitals in the highest principal shell first. Thus say we want singly ionized zinc, the electron configuration would be: Zn + : 1s 2s 2p 2p 2p 3s 3p 3p 3p 4s 3d 3d 3d 3d 3d Zn + : 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 which now would be paramagnetic, due to the one unpaired electron in the 4s orbital. It should be noted that Zn + is extremely rare while most compounds contain Zn 2+. In this lab the student will be given several known compounds and be asked to determine whether they are diamagnetic or paramagnetic. Furthermore, the student will be asked to write the electron configuration of the elements in the sample and compare these configurations to their experimental results. Other than magnetic properties, the electron configuration of these compounds determines most of their properties, including reactivity, atomic and ionic radii, boiling and melting point, and color. Acknowledgement: This experiment is based on A General Chemistry Laboratory Experiment Relating Electron Configuration and Magnetic Behavior, by Wayne H. Pearson, J. Chem. Educ. 2014, 91, II. PROCEDURE Saf ety P recautions: Several of the samples, though sealed, are toxic and carcinogenic. DO NOT OPEN THE VIALS! Take care with the vials. The vials may shatter if dropped, thus causing their contents to escape. Five points will be deducted from your lab for each vial you break. The magnets are stronger than those the reader may be use to. Take care not to pinch your fingers when placing the magnets on the balance. In this lab and in every lab students must wear goggles. The students, working in pairs, will obtain ten sealed vials. DO NOT OPEN THE VIALS! Having the vials preloaded saves chemicals, time, and minimizes the possibility of exposure. These vials are labeled with their contents. Nine of the vials will contain knowns, while the last vial will be an unknown. This unknown will be one of the previous nine. Zero the analytical balance, then place the magnet on the balance and record its apparent mass; the magnet can be found next to the balance. The term apparent mass is used since there will be some interaction between the magnet and the balance altering the measurement. Record this value in the data sheet. Elevate the vial containing the sample approximately 0.5 cm above the magnet. Do not allow the vial to touch the balance or the magnet for this will alter the reading. Record the newly displayed mass in the data sheet. If the apparent mass has decreased in the presence of the sample, then the sample is attracting

4 the magnet, and the sample is paramagnetic. If the apparent mass has increased in the presence of the sample then the sample is being repelled by the magnet; consequentially the magnet is also being repelled, and the sample is diamagnetic. Perform two measurements then return to your lab bench and calculate the number of unpaired electrons for those two samples based on the oxidation numbers of the elements. Repeat until all samples have been measured. 4 III. PRE-LAB QUESTIONS 1. (2 points) Below you are presented with sets of quantum numbers: (n, l, m l, m s ). Only one is a valid set. Which is it? A) (4,3,2,1) B) (3,3,2,1/2) C) (3,1,-1,-1/2) D) (2,0,0,0) E) (5,3,-4,1/2) 2. (2 points) What is the maximum number of electrons can a d subshell can hold? Please enter an integer. 3. (2 points) One of these subshells is not possible, which is it? A) 3d B) 4d C) 1p D) 4s E) 2s 4. (2 points) Which of the following cations would you expect to be diamagnetic? A. Cu 2+ B. Zn + C. Al 2+ D. Ti 4+ E. Fe (2 points) Which of the following anions would you expect to be paramagnetic? A. As 2 B. S 2 C. P 3 D. N 3 E. Si 4 6. (2 points) How many unpaired electrons would you expect on aluminum in aluminum (III) oxide. Enter an integer.

5 5 7. (2 points) How many unpaired electrons would you expect on iron in iron II sulfide. Enter an integer. 8. (2 points) How many unpaired electrons would you expect on Vanadium in V 2 O 3 Enter an integer. 9. (2 points) How many unpaired electrons would you expect on iron in [Fe(H 2 O) 6 ] 3+ Enter an integer. 10. (6 points) Challenge question: This question is worth 6 points. As you saw in problem 9 we can have species bound to a central metal ion. These species are called ligands. In the past we have assumed all the d orbitals in some species are degenerate; however, they often are not. Sometimes the ligands bound to a central metal cation can split the d orbitals. That is, some of the d orbitals will be at a lower energy state than others. Ligands that have the ability to cause this splitting are called strong field ligands, CN is an example of these. If this splitting in the d orbitals is great enough electrons will fill low lying orbitals, pairing with other electrons in a given orbital, before filling higher energy orbitals. In question 7 we had Fe 2+, furthermore we found that there were a certain number (non-zero) of unpaired electrons. Consider now Fe(CN) 6 4 : here we also have Fe 2+, but in this case all the electrons are paired, yielding a diamagnetic species. How can you explain this? A) All the d orbitals are degenerate. B) There is 1 low lying d orbital, which will be filled with two electrons before filling the 4, assumed to be degenerate, higher energy orbitals. C) There are 2 low lying d orbitals, which will be filled with 4 electrons before filling the 3, assumed to be degenerate, higher energy orbitals. D) There are 3 low lying d orbitals, which will be filled with 6 electrons before filling the 2, assumed to be degenerate, higher energy orbitals. E) There are 4 low lying d orbitals, which will be filled with 8 electrons before filling the 1 higher energy orbital.

6 6 IV. POST-LAB QUESTIONS 1. Chromium, Cr, has 24 electrons. Write out the entire electron configuration for chromium using spdf notation. 2. How many unpaired electrons would you expect for chromium in [Cr(H 2 O) 6 ] 3+? Is this a paramagnetic or diamagnetic material? 3. Cobalt, Co, has 27 electrons. Write out the entire electron configuration for cobalt using spdf notation. 4. How many unpaired electrons would you expect for manganese in KMnO 4? Is this a paramagnetic or diamagnetic material? 5. This same experiment is performed on the international space station. What is the primary issue with performing this experiment in the absence of gravity? Design an experiment to compensate for this. As always, you do have duct tape.

7 7 V. CHEMICAL HAZARD AWARENESS FORM: MAGNETIC BEHAVIOR AND ELECTRON CONFIGURATION OF COMPOUNDS This form is to be turned in prior to beginning your experiment. Name: Section: RUID: Experiment: List below any chemicals used in this experiment. Also list their hazards and any handling precautions for these chemicals.

8 8 VI. DATA SHEET: MAGNETIC BEHAVIOR AND ELECTRON CONFIGURATION OF COMPOUNDS Name: Section: RUID: Instructor: Please complete Table II. The mass of the magnet is the original mass measured before suspending the compound above the magnet. The apparent mass is the mass of the magnet during suspension of the compound. The change in mass is the difference between the previous two masses. If the change in mass is a positive value then the sample is diamagnetic. If the change in mass is negative then the sample is paramagnetic. Record all masses in grams. TABLE II: Data # Compound mass of magnet [g] apparent mass [g] Change in Mass [g] 1 KCl 2 MgO 3 MnO 2 4 CaCl 2 5 Mn 2O 3 6 CuI 7 CuO 8 Fe 2O 3 9 MgBr 2 10 Unknown

9 9 VII. RESULTS: MAGNETIC BEHAVIOR AND ELECTRON CONFIGURATION OF COMPOUNDS Name: Section: RUID: Instructor: Results: Complete Table III by indicating the number of unpaired electrons: n and whether the compound is paramagnetic: P or diamagnetic: D. TABLE III: Results # Compound n P or D 1 KCl 2 MgO 3 MnO 2 4 CaCl 2 5 Mn 2O 3 6 CuI 7 CuO 8 Fe 2O 3 9 MgBr 2 10 Unknown N.A. Conclusions and Comments: What compound is your unknown?