Unit 6: Introduction to Chemical Bonding
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Unit 6: Introduction to Chemical Bonding |
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Introduction:
Unit 6 " Introduction to Chemical Bonding" is the second of four Units focused on atomic structure and chemical bonding. Chapter 8 in the text and the related "screens" on the ROM, extend atomic structure concepts to include the fourth quantum number, m(s) (aka as "m-sub s"), electron spin. The values of m(s) are either +1/2 or -1/2 , often symbolized as an "up arrow" or a "down arrow". Two additional "rules" associated with quantum numbers go under the banners of Pauli Exclusion Principle which states that no two electrons in the same atom can have the same four quantum numbers, n, l, m(l) and m(s) and Hund's Rule of Maximum Spin. Hund's rule says if more than one spin state is available with equivalent energy, choose the state with the most "up-arrows".There are several symbolic systems introduced in the unit which are guided by thePauli Principle and Hund's Rule and used to characterize the electronic structure of atoms and ions . You should study this unit until you are completely familiar with the symbols for electronic configurations of atoms and ions.
The "n+l" rule for determining the relative energy of the various sub-shells (atomic orbitals ) in a atom is the last rule associated with quantum numbers you will need to know. Simply stated, if the sum of "n", the principle quantum number and "l" are the same, fill the atomic orbitals with the lower n value first. For example, n = 3 and l = 0 (3s) has the same n+l as n = 2 and l = 1 (2p) both are n+l = 3 so the "rule" guides you to complete the 2p orbitals before filling the 3s. You will see more on this in the unit.
After examining a few of the periodic trends for ionization energy and ionic size, Unit 6 introduces the basics of covalent bonding in chemistry, the "Lewis Dot Symbols" for the atoms in second row of the periodic table, particularly, carbon, nitrogen and oxygen.
Unit Objectives: After completeing Unit 6, you should:
1. Be familiar with the fourth quantum number, m(s) electron spin. Give the assigned quantumn numbers n, l, m(l) and m(s) for electrons and ions up to calcium, element # 20 using the Pauli Exclusion Principle.
2. Know how to apply the "n+l" rule and be able to reproduce electron configurations using spectroscopic notation 1s, 2s, 2p, 3s and so on, for elements and ions up to calcium (#20) on the periodic table.
3 . Be able to use the "box diagrams" for s and the three p atomic orbitals to demonstrate the effect of Hund's Rule on the electron configurations of the second and thied row elements.
4. Be able to use the periodic table to organize the "perodic trends" of the Representative Elements, (Group A Elements) for atomic size, ionic size, ionization energy and electron affinity.
5. Be able to explain the difference between ionic interactions and covalent bonding.
6. Be able to give the Lewis Dot Symbols for the elements H, C, N, O and the halogens and show how to draw Lewis Dot Structures for some simple, single covalent bonded examples.
Reading and Study Assignment:
Text: Chapter 8 ,"Atomic ElectronConfigurations and Chemical Periodicity" pages, 357 to372 and 376 to 386. Omit "Electron Configurations for the Transition Metals" pages 373 to 375 and Section 8.7 "Chemical Reactions and Periodic Properties", pages 386 to 390.
Chapter 9, "Bonding and Molecular Structure" pages 398 to 407.
CD-ROM: Chapter 8 and the first few "screens" in Chapter 9.
Electron Spin and the Pauli Exclusion Principle:
This diagram, from the "Chemistry" by Atkins and Jones, if it is "clear" and readable for you, represents a summary of the quantum numbers, n, l and m(l) and spectroscopic symbols from Unit 5. The first section in the text and Screens 8.2 and 8.3 introduce the last quantum number, m(s) for electron spin, either "spin-up" or "spin-down", +1/2 or -1/2. The combination of all four quantum numbers are needed to complete the "address" of any particular electron in an atom.
The Pauli Exclusion Principle, discussed in Section 8.2, page 360 of the text and Screen 8.4 of the ROM states, "no two electrons in the same atom can have the same set of four quantum numbers, n, l, m(l) and m(s)". This principle imposes a limit on the total number of electrons that can occupy a particular atomic orbital set. For example, only two electrons maximum can be placed in any orbital given by n. l. and m(l). See Table 8.1 on page 362 for a complete summary. You should be able to reproduce this sort of table for the n =1 to n = 4 energy levels with the associated l, m(l) and m(s) values.
Workbook Notes:
Omit the workbook questions for Screens 8.2, 8.3 and 8.4 However be aware that a "spinning" electron is portrayed as a little magnet, therefore a single spinning, "unpaired" electron would be deflected by a magnetic field. (Screen8.2) This concept of deflection of a unpaired electrons in a magnetic field is the basis for an analytic instrument used in chemistry called "electron paramagnetic resonance spectrometer," epr-spectroscopy.
For Screen 8.4, complete the statement of the Pauli Principle and give a set of four quatnum number that could represent an electron in a 2s orbital , a 3s orbital, a 2p orbital and a 3p orbital. For example a 3p orbital could be "addressed" as 3, 1, 0, +1/2 . You come up with another address for a 3p electron. There are five other addresses for "3p" that will obey the Pauli Principle.
Atomic Subshell Energies:
Understanding the order of subshell energies is the most important concept in this unit, the "n+l Rule". Subshell energies creates the filling scheme for electron configurations in atoms and ions. See Text page 365 and Screen 8.5. In addtion you should practice the questions and problems in the section since the symbolism introduced here is used to discuss covalent bonding. For most situations, it will be sufficient to be familiar with the notation up to n = 4.
Answer the question in the workbook for Screen 8.5 and do Workbook Exercise 8.1 these are n+l rule applications. Check with the text Exercise 8.1 on page 366.
Electron Configurations
Read through Section 8.4 in the text, beginning on page 367 and view Screens 8.7 and 8.8. The "box" and spectroscopic symbols for electron configurations is the culmination of the introduction of quantum theory in a general chemistry course. You will use the symbols in future discussions about chemical bonding, so you must become familiar with them now. The "box" diagrams are the best for showing the effect of Hund's Rule on electron configurations. I have included a set of box diagrams from the text by Atkins and Jones. Similar diagrams are in your text on page 369
Note: This diagram shows the "unpairing" of the 2p electrons as they fill the three boxes, 2p(x) ,2p(y) and 2p(z) as you go from boron to nitrogen. Oxygen, with 6 outershell electron in the n= 2 level, is forced to pair electron spin, as shown above. I don't think the spectroscopic notation will show, but I wanted you to see the "box" scheme for these elements. I think it will be OK. Also, check how the Lewis Dot symbols correspond to the number of "outershell" electrons in each case. You will practice writing a lot of these 1s,2s, 2p etc. configurations.
Workbook Notes :
Screen 8.7 is an introduction to "box" notation, "Noble gas" notation and "spectroscopic" notation. On the "sub-screen" Atomic Electron Configuration click on and study the items under "Periodic Blocks", "Hund's Rule" omit "Trensition Metals" and "Anomalies". Concentrate on the "s" and "p" block elements, these are called the Group A or Main Group Elements. The periodic properties of these elements follow predictable trends. Also, the non-metals plus H, hydrogen are the principle elements involved in covalent bonding. Make sure you understand all three symbolic presentations of electron configurations. For Screen 8.7, you should complete workbook questions 1 and 2. Replace "vanadium" with potassium in # 3 and "cerium" with chlorine in # 4 . and complete these exercises. Omit # 5. Finally, complete the suggested activity for Hund's Rule in #6 for the third row elements Al, Si, P and S. You can check #6 against the second row elements B, C, N and O, the have similiar configurations only using "2p" electrons instead of "3p" ones.
Omit Exercise 8.2 in the workbook and substitute Exercise 8.2 in the text , page 372. Check answer in Appendix, page A-41.
Screen 8.8 Remember ion formation invloves the removal of electrons to form cations and the addition of electrons to form anions. In order to give the electron configuration for cations or anions, you count the total number of electrons in the ion and "fill" the energy levels using the same rules as neutral atoms. Again, we will focus on the s and p block elements for these exercises. For Screen 8.8 complete questions 1 and 2 . Note both of these exercise are for isoelectronic configurations, that is the configuration K(+) or Ca (2+ ) is [Ar], the noble gas argon configutration where all species have 18 electrons. Similarly, the configuration for O(2-) or F(-) is [Ne], the noble gas neon configuration with 10 electrons. Omit Exercise 8.3
Periodic Trends:
You should remember a few general principles for correlating these periodic trends. 1) Positive ions, cations, are formed by removing one of more electrons form the neutral atom, the "effective"nuclear charge increases for the remaining electrons in the positvie ion thus, the ionic radius is now smaller than the neutral atom radius. See the bottom of page 384 comparing Li and Li(+). In the case of negative ion formation, anions, electrons are added to the outershell of the neutral atom, the effective nuclear charge is less per electron thus, the ionic radius is larger than the neutral atom radius. See text page 385 comparing F and F(-). Know the definition of ionization energy (page 379) and electron affinity (page 381). The terms "exothermic" energy releasing and "endothermic" energy absorbing can be use to describe the enegy flow in an electron affinity event or in the case of ionization.
For electron affinty, an electron is attracted to a nectral atom to form a negative ion , in chemistry, an "attraction" event is energy releasing, that is the resulting specie, a neagative ion int his case, is in a "lower" or more favorable enegy state than either the neurtral atom or the electron initially. Such a "reaction" could be viewed as;
F(g) + electron ---> F (-) ion + energy
Since electron affinity results in an "energy release" relative to the "reactants", electron affinties are listed as negative numbers, See Table 8.5 , page 381. This can be viewed as being analogous to electron emission, as the "free electron" is trapped by the fluorine, it "fall' into lower enegy states, releases energy and is "captured" by the attractive force of the fluorine nucleus.
The ionization of an atom to a positive ion requires the reomval of a electron for the attractive forces of a nuclear charge. To remove an electron requires an energy input, and is an endothermic event,
Na (g) + energy ( "first ionization energy") ---> Na(+) gaseous cation
The ionization energies listed in Figure 8.13, page 380 are positive numbers because this reaction requires an absorption of energy and is called an endothermic reactive event.
You should know general trends in the quanitities, atomic and comparable ionic radii, ionization energy and electron affinity as you move left to right across the period or "row" of the Group A elements and as you move from top to bottom down a Group A Family = (column).
Workbook Notes
For Screen 8.9, complete question 1 in the workbook.
For Screen 8.10, comnplete questions 1, 2, and 4 . In the case of question 4, pick one or two examples. Note the common distance at the "atomic' level is the picometer = 10E-12 meters, and is symbolized as "pm". So an atomic radius of 77 pm for a carbon atom means the C atom radius is 77x10E-12 meters. Do workbook exercise 8.4 and check against Figure 8.10, page 377.
For Screen 8.11, complete questions 1 and 2.
For Screen 8.12 , complete questions 1, 2 and 3
For Screen 8.13 (Last Screen assigned in Unit 6), complete 1, 2 and 3. Exercise 8.5 in the workbook is the smae as Exercise 8.7 in the text, page 383. Do it and check A-41 for answers.
Valence Electrons and Lewis Dot Strucutres:
The last concept in this unit, is to identify the number of valence or outershell electrons in the Group A elements and use these electrons to "create" electron pair or covalent bonds. G.N. Lewis described covalent bonds as "shared electron pair" bonds in the 1920's. His ideas, although refined since he first formulated them, remain as the primary comcept for bond formation in compounds involving the non-metals and hydrogen. The diagram in the Hund's rule section shows the Lewis Dot Symbols for B, C, N and O. More symbols are shown in Table 9.1 page 399. In order to begin to make Lewis structures for simple, covalent compounds using the second row elements, the halogens and hydrogen, you need know the dot symbols for the elements and the "octet rule". The Lewis Octet Rule states that the second row elements, C,N,O and F, form covalent compounds by completing their valence shell (outershell) by sharing electrons with other atoms until there a total of eight in this shell. For example C with four valence shell electrons will form four electron pair bonds so it will have eight total electrons in its valence shell. N will form three bonds, O will form two bonds and F needs only one, electron pair bond to complete its "octet".
Screen 9.2 complete questions 1 and 2
Screen 9.3 omit the questions in the workbook but you should think about the difference between an ionic compound like NaCl, sodium chloride, and a covalent compound like CH4 (g).
Screen 9.4 complete questions 1, 2 and 3.
Screen 9.5 complete questions 1 and 2. Don't worry if you don't understand all the strategy for drawing Lewis Structures.We will begin Unit 7 with "Screen 9.5"
e-mail Task (15) Due by Friday evening 2/27
a) What do the diagrams in this Unit look like on your screen and or print out. Are they ledgible?
b) Explain why the first ionization energy of an element generally increases as you move across a row of the s and p block elements.
Study Assignment:
Workbook Chapter 8 - 1 (13), 5 (25), 8 (39), 9 (43), 10 (49), 11 (55), 12 (57), 13 (59), 17 (75), 18 (80) and 20(94). No problems from Chapter 9 this time.
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