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Introduction: The next four units in Chem 211 will explore the current view of atomic structure and chemical bonding. Atomic structure in its present form, emerges from the advances in physics in the first few decades of this century. Chemical bonding, the mutual interaction of atoms with each other, is described from the perspective of electrons in atoms, the chemically relevant part of the theory of atomic structure. If you asked Joe Chemist for a definition of "chemistry". He might suggest that,"Chemistry is the study of what happens to the outer shell electrons". Keep this comment in mind, you are about to begin a study of atomic structure focused on identifying which electrons in an atom are the "outer shell ones". After completing Unit 5, you will are able to give the electronic configuration of an atom including the number and type of outer shell electrons. The next unit will introduce the interaction of these electrons between atoms resulting in "chemical bonding".

The format of Unit 5 will be changed from the previous four units. Since the discussion of chemical bonding is more qualitative and less quantitative than stiochoimetry, these units will attempt to "guide" you through the text, the CD- ROM and the workbook. The quiz and e-mail activities will remain but there will be no FAQ page. Each unit will be constructed with more extensive direction for reading the text and using of the CD as a learning tool. Since certain features of the CD and associated questions in the workbook can function well for introducing the concepts of atomic structure, I plan to try to develop units to "guide" you through these activities more than before. Please let me know if this scheme is helpful. Thanks ---MD

Unit Objectives: After completing Unit 5 , you should be able to:

1. define the terms wavelength, frequency and speed of light and write an expression relating these to each other. [Frequency (sec -1) = Speed of light, c (m/sec) / wavelength (m)]

2. calculate "photon" energies using Planck's expression., e= hv or E= N hv . You must be able to calculate both the energy per particle and the energy per mole of particles, given the appropriate wavelength or frequency.

3. describe the Rutherford-Bohr "planetary" model or "shell" of the atom.

4. explain qualitatively what the term quantized energy levels means for electron energy states in the atom as postulated by Niels Bohr.

5. explain how electrons move from one energy level to another for "hydrogen like" energy levels when energy is absorbed (absorption spectra)or emitted (emission spectra) by an atom. ( Can any amount of energy cause the absorption or emission by electrons in atoms?)

6. understand that an atomic orbital represents a region of space where there is a high probability of find one or two electrons of the particular "subshell" represented by the atomic orbital "picture"

7. describe the energy level and atomic orbital representation for the quantum numbers, n, l and m(l). Reproduce the atomic orbital representations for an "s" orbital ( l = 0 ), the three " p" orbitals (l = 1 ), giving their shapes and geometric directions. You should be familiar with the shapes of the d orbitals , ( l = 2 ) but you will not be expected to reproduce them. [ See pages 343 and 346 and CD Screen 7. 13 ]

Reading and Study Assignment:

Text: Chapter 7, "Atomic Structure" pages 312 to 349.

CD-ROM: Chapter 7

Unit 5 "Guide": A quest for the quantum numbers "n", "l" and "m(l)" aka "m - sub l"

Screen 7.2 shows a historical time line for the discoveries leading to the current theory of atomic structure. The "time line" extends from about 1860 to 1930 ish which is the important period for the development of atomic theory in Europe and the United States. You can "move" around the ROM using this "line" or the "screen targets"at the bottom of the screen. Modern atomic theory is generally a twentieth century discovery.

Electromagnetic Radiation: Section 7.1 and Screens 7.3 and 7.4

This section introduces the terms, wavelength [Greek letter lambda], frequency [Greek letter nu] and speed of light, c. You should discover by looking at the "frequency related to wavelength" demonstration for the visible part of the electromagnetic spectrun, that the frequency is the reciprocal of wavelength multiplied by the speed of light. Or, as shown in the text, wavelength times frequency is a constant, the speed of light. It is hard for me to "write" this equation but it is shown in both the ROM and the text ( equation 7.1, page 315). For purposes of calculating frequency from a given wavelength, you can use the speed of light to be 3 x 10 E+8 meters/sec not 2.998 m/sec. You must be careful with units in these calculations. For example, the wavelenghts common to "electrons in atoms" are very short and given in nanometers ( 10E-9 meters) and must be changed to meters before multiplying by the speed of light if you are trying to calculate the frequency. (See example in the middle of page 315).

Since wavelength and frequency of the electromagnetic spectrum are related as reciprocals, this means as the wavelength gets shorter the frequency goes higher and conversely, as the wavelength gets longer the frequency gets smaller. In the next section, Planck shows that the electromagnetic energy in joules/ photon or joules/(mole of photons) is directly related to frequency. Thus the higher the frequency the higher the relative energy of a "photon" of electomagnetic radiation.

Screen 7.4 [Figure 7.3 page 315] shows the "visible" portion of the "e-mag spectrum" It is a small portion of a spectrum that spans frequencies from 10 hertz or "cycles per section" to 10 x10E+24 hertz. Visible light is in the region given by wavelenghts of 700 nm "red end" to 400 nm "blue end" Note on Screen 7.4 how the shape of the wave changes as you move through the visible spectrum.

Workbook Notes:

Screen 7.3

Do all the questions. For question 2, consult figures 7.2 and 7.4 and Exercise 7.2, page 318. Additional things to think about regarding your "standing wave pictures" How many "nodes" points where the wave crosses the axis, exist for your 2 cm wavelength drawing. Refer to the Exercise in the text,page 315 and answer the questions related to the same idea.

Screen 7.4

Answer the question in the workbook for this screen. The principle objective of the exercises is to show that frequency and wavelength are inversely related to each other. The answers to questions 2,3 and 4 can be found answers to Exercise 7.1 in the Appendix page A- 40. Also consider,

What has higher frequency, microwaves or UV (ultra-violet) waves.?

If energy is a direct function of frequency, which of the following has higher energy ?

a) blue light, 400 nm or red light 700 nm

b) FM radio waves, about 3 meters or microwaves, 0.03 meters

Energy and Photons: Section 7.2 and Screen 7.5

The results of Planck's radiation experiments is the equation relating frequency of some radiation, "packet of photons" and energy. The constant in this equation bears Planck's name and is symbolized by "h" Planck's constant is 6.62 x 10E-34 joule-sac/ photon. and the equation is shown on the screen or page 319 of the text. See example 7.2 on page 323 for calculating photon energies using Planck's equation. The difference between individual photon energy at some given wavelength or frequency and the energy of a "mole " of photons is Avagrodro's number, N where N = 6.02 10E+23 photons/ mole in this case.

Workbook Notes:

Screen 7.4

Answer questions 1 and 2 . Instead of doing the "Exercise 7.1 " in the workbook do the similar exercise 7.3 in the text on page 323 and check your answers to it.

Atomic Line Spectra - Bohr Model: Section 7.3 and Screens 7.6 and 7.7

The Rutherford experiments established the subatomic nature of the nucleus, as a dense, positively charged "core" of the atom. The question Bohr explored was how could electrons, negatively charged particles be expected to stay in stable "orbits" around the nuclear core ? Wouldn't the electrons be attracted to the nuclear core and simply "fall" into resulting in the collapse of the atom? To answer the question of why the atoms didn't collapse, Bohr postulated the "quantized"electron , a species allowed to exist in only certain energy levels. This postulate is the working model for the Lewis Bonding Theory in Chemistry . Atomic line spectra stand as experimental "proof" of the existence of only certain energy levels for electrons in atoms as suggested by Bohr. A good example of "line spectra" is Figure 7.10, page 326. The is what Screen 7.6 is demonstrating as an experiment using a light source involving an electrically excited, pure gas such as neon, then passing the photons through a prism to separate the photons into their respective energy pattern.

Workbook Notes :

Screen 7.6

Answer questions 1 and 2 omit 3 and 4 Review the discussion of the term quantized energy levels in this section in the text . Why did Bohr make this assumption? How does the existence of atomic line spectra support the concept that electron energy in atoms is "quantized"? The existence of these"Bohr hydrogen like "energy levels" is an important part of the quest for "n" "l" and "m(l)" The energy level is given the principle quantum number, n and n can have values beginning with 1 up to 7 or so all the atoms known to us at this time. The really important values of n (energy level or "shell" ) are 1,2,3,4,5,6 and 7 for chemists.

Screen 7.7

Answer these questions: Omit the questions in the Workbook for this screen.

a) What are the assumption of the Bohr Theory of the atom?

b) What does it mean to say electrons can only occupy certain energy levels in an atom?

c) What is the value of n for the " ground state" of a hydrogen atom? Remember a hydrogen atom of atomic mass = 1.0 , consists of a proton and an electron. When you refer to the Bohr "hydrogen-like " model this is the nature of that species, 1 proton and 1 electron.

Wavelike Character of an Electron: Section 7.4 and 7.5 and Screens 7.8,7.9 and 7.10

This section completes our "quest" for the quantum numbers n,l and m(l) as viewed by chemists. The quantum theory behind the "wavelike" character of an electron is "beyond the scope of this course, covered in more detail in Chem 352. However, what you need to know is electrons are viewed either as "particles" or as "wave packets" as is convenient for so theoretical treatment, The results of the quantum theory briefly discussed in the text, introduces you to the "l" or angular momentum quantum number and the "m(l)" [ read "m-sub l"] quantum number.See Table 7.1 for a summary of the allowed values for these quantum numbers. Be aware, the values of l depend on the value of n and the values of m(l) depend on the value of l. You should memorize these possible values up through n=4 for this course. However once ;you understand the way the quantum numbers are related to each other, you could assign values up to n= 8 and beyond, much farther that necessary in chemistry.

Table 7.1 page 340 can be summarized as (2/17 Note: Original had error in last box, l = 3 "d" orbitals, now correct ed to read ," l = 2 "d" orbitals. If l =3, the "f" orbitals with m(l) = -3,-2,-1,0,+1,+2,+3 are generated.)

n = 1 ( lowest energy level "closest to nucleus") l = 0 m(l) = 0

n = 2 ( next higher energy level)	l = 0 "s" orbital	m(l) =0
or for same n value	l = 1 "p"orbitals	m(l) = -1,0 +1

n = 3 (next higher energy level)	l = 0 "s" orbital	m(l) = 0
or for same n value	l = 1"p"orbitals	m(l) = -1,0,+1
or for same n value	l = 2"d"orbials	m(l) = -2,-1,0,+1,+2

You should complete a "box" for n=4 and check it against Table 7.1.

Workbook Notes:

Screen 7.8

Look at DeBroglie's Equation and answer question 1 (a) and (b) Omit 2.

Screen 7.9

No questions Note the Heisenberg Uncertainty Principle is the basis for the probability of location of an electron in some spatial region. The theme of "atomic orbital diagrams" is that these"picture" represent the spatial region where you would have a high probability of finding an electron because its quantum numbers, n, l and m(l) give that spatial orientation.

Screen 7.10

View this screen to complete the "historical" path to the modern theory of the atom. Omit all the workbook questions for this screen except this variation

.How many quantum numbers must be used to describe an atomic orbital. What does each quantum number represent? And how are their values assigned relative to each other?

Shapes of Atomic Orbitals: Section 7.6 and Screen 7.12 and 7.13 ( Note You can view screen 7.11 but it seems confusing to me and not needed.)

This is the important stuff to know for the next unit and is the end of the "Quest" You know the origins of the quantum theory of the atom and of the three quantum numbers n, l and m(l) used to describe the "location " of electrons in atoms. The quantum number n, principle quantum number is indicative of the energy level of an electron and controls the possible values of l. The quantum number "l" dictated the "shape" of the atomic orbital and is designated as an "s" shape for l = 0, a "dumb bell" shape identified as "p" atomic orbital for l = 1 and finally., for Chem 211 purposes, l = 3 gives shapes associated with the atomic orbital set called "d" orbitals. The orbital angular momentum quantum number, m(l) gives the geometric direction of the "l" shapes. (See page 343 and 346 for example or Screen 7.13.).

Workbook Notes:

Screen 7.12

Complete all the questions for this screen. You might find Table 7.1 and Fig 7.17 in the text as useful references for the questions. Exercise 7.2 in the workbook is the same as Exercise 7.7 , page 341 in the text, do it and check your answers.

Screen 7.13 Omit the workbook question for this screen. Instead make a sketch of the following atomic orbitals correctly labeled axis in the case of the 2p orbitals. Sketch a 2s, 2px, 2py and a 2pz orbital and check your sketches against those in Figure 7.17 page 343.

e-mail Activity: Due by Friday evening, 2/20

See if you can find the word/s, "quantized" or "quantum theory" in a dictionary. If so, how are they defined. In the unit, you have discovered Bohr had to postulate "quantized energy " levels for electrons in atoms in order to keep them in "stable" orbits around the nucleus. Briefly explain how the existence of line spectra supports the quantized nature of the electrons in atoms.

Study Questions for Unit 5:

Workbook "Study Questions" 1 (19),2 (23), 5 (30), 7 (39), 8 (46), 9 (47), 10 (49), 11 (51), 12 (53), 13 (79) and 16 (45).

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