(3) Model: a nucleus of ( ) charge that also contains ______________ nucleus is encircled by e-\'s located in definite orbits (or paths). e-\'s have ___________ energies in these orbits e-\'s do not lose energy as they orbit the nucleus(4) Mechanical Model ( Wave Mechanical Model) no definite ____________ to the e- path (?fuzzy\" cloud) orbits of e-\'s based on the _________________ of finding the e- in the particular orbital shape. . - PowerPoint PPT Presentation
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1. Ch. 11 Notes---Modern Atomic Theory
3. Bohr Atomic Model
5. Quantum Mechanical Model
6. Quantum Mechanical Model
9. Figure 11.15: The difference between continuous and quantized energy levels.
10. Quantum Numbers Describe the ______________ of the e-’s around the nucleus.
Quantum #’s are sort of like a home _____________ for the electron.
This information about the location of the e-’s in an atom can be used to:
(1) determine chemical & physical _____________ for the elements.
(2) show how the _______________ __________ is organized.
(3) show _____ and _____ elements combine to form compounds.
11. The Four Quantum Numbers Principal Q. #: Describes the _____________ that the electron is from the nucleus. The bigger the number, the ___________ away the electron is.
Example: (1=closest, 2, 3, 4...farther away)
These distances are sometimes called _______________
12. Orbital Q. #: Describes the __________ of the electron’s path around the nucleus with a letter: (s, p, d, & f) These are sometimes called “_____________”.
s=_____________ cloud; p=_____________ or a 3-D figure 8;
13. d & f orbital shapes are complex ________- _______________ ellipsoids, and some d’s and f’s are an ellipsoid with a doughnut or two around the middle.
All of these orbital shapes are based on the probability of finding the electron in the cloud.
14. Figure 11.22: How principal energy levels can be divided into sublevels.
15. A Way to Visualize s, p, d, and f sublevels
16. Magnetic Q. #: tells how many _________________ in 3-D there are about the nucleus for each orbital shape.
s=___ orientation p= ___ orientations... (x, y, and z)
d= ___ orientations f= ___ orientations
The orientations are represented with a line or a box.
Examples: ___ This means a __________ orbital at a distance of 1s “__” (close) to the nucleus. This orbital is centered about the x, y, and z axis.
? ? ? This represents an ___________ orbital with its 4p ____ possible orientations at a distance of “___” from the nucleus.
17. Figure 11.23: Principal energy level 2 shown divided into the 2s and 2p sublevels
18. Spin Q. #: describes how the electron in an orientation is spinning around the nucleus. This spin can be thought of as “____” or “________”. (Some like to imagine it spinning “clockwise” and “counterclockwise”.) The spin is represented as an ___________ in the direction of the spin.
Example: ? This represents one electron in a _________ 2s orbital with spin “____” at a distance of “___” from the nucleus.
Remember, the four quantum numbers tell us the location, or “address” of each electron in an atom.
This information is vital in understanding the layout of the Periodic Table and the reasoning behind why and how atoms form bonds.
19. Electron configurations are notations that represent the four Quantum #’s for all of the electrons in a particular atom. Here are the rules for these notations:
Rule #1 (Aufbau Principle): Electrons fill ________ energy orbitals first.
Examples: 1s would be filled before ____
3s would fill before ____
21. Rule #2: Only ___ electrons can fit into each orientation.
Example: ___ ___ not ____
1s 2s 1s
Rule #3 (Pauli Exclusion Principle): Electrons in the same orientation have ______________ spin.
Example: ___ not ___
Rule #4 (Hund’s Rule): “_______________ rule”---> Every “?” in an orbital shape gets an electron before any orientation gets a second e-.
Example: ??? not ???
22. Rule #5: The Exceptions
Half-filled or completely filled d & f sublevels have ________ energies and are more stable than partially filled d’s and f’s.
This means that an atom can “borrow” one of its “s” electrons from the previous orbital to become more stable.
Example: ___ ___ ___ ___ ___ ___
becomes ___ ___ ___ ___ ___ ___
Because the 4d sublevel is now full, the atom is at a ________ energy state and therefore _________ stable.
23. Electron Configurations Practice Problems:
Write the electron configuration notation for each of the following
Shorthand Method: Assumes we already know about the # of ?.
24. How Electron Configurations Relate to the Organization of the Periodic Table
25. Figure 11.31: Orbitals being filled for elements in various parts of the periodic table.
26. Electron Configurations & Properties How do electron configurations relate to the chemical and physical properties of an element?
All elements with the _________ outer shell e- configurations have ________ properties.
This means that elements in the same ____________ group have similar properties.
Examples: (1) Li, Na, K, Rb, and Cs all have __ lone “__” e- for their last orbital... (_____, _____, _____, etc.) This makes all of them ___________ reactive. They all react with __________ to produce hydrogen gas.
(2) Ne, Ar, Kr, Xe, and Rn all have the outer energy level completely __________ with electrons...(________, ________, ________, etc.) This makes all of them ______________. They do not produce __________________!
27. More Practice Problems
(1) Which element has its last electron as a 4p5? ___________
(2) Which elements are similar in properties as Bromine? __________
(3) What would the outer shell electron configuration look like for the element underneath Radon, (Rn)?
(4) Which electron is added after 6s2? ________
(5) Which element would “borrow” a 5s electron to get a half-filled “d” sublevel? ___________
(6) What is the shape of the last orbital filled in Calcium, (Ca)? _____
(7) How many electrons are in the last “p-orbital” of Sulfur, (S)? ____
28. Electromagnetic Radiation Any wave of energy traveling at a speed of ___________ is called electromagnetic radiation.
There are many types of electromagnetic radiation and each type has a different _______________ and _______________.
Here are the types of electromagnetic radiation from longest to shortest wave or lowest to highest frequency. These are also in order from lowest to highest energy.
29. Electromagnetic Radiation
30. Electromagnetic Radiation (1) Radio Waves -- used in __________________
(2) Microwaves-- broadcasts TV signals and used to _____ _______.
(3) Infrared (IR) -- we feel this as _____; _________ & ______ can “see” this.
31. Infrared Vision
32. Electromagnetic Radiation (4) Visible Light -- the only radiation we can detect with our eyes. It can be separated into the colors of the spectrum with a __________.
(5) Ultraviolet (UV) -- gives you a _____________; _________ can “see” this; some of this radiation from the sun gets blocked by the ___________ layer
33. Electromagnetic Radiation (6) X-rays -- used in medicine
34. Electromagnetic Radiation (7) Gamma Rays-- some radioactive substances give it off
_______________Rays – These are not part of the EM spectrum… They are high energy particles (mostly protons); They cause the northern lights.
35. How Light is Produced When atoms get hit with energy (by _____________ them with electricity or by ____________ them up), the electrons absorb this energy and __________ to a higher energy level. Figure (a)
As they immediately fall back down to the “____________ state”, they give off this energy in the form of a particle of ___________ (or other types of electromagnetic radiation) called a _____________. Figure (b)
36. How Light is Produced Each photon emitted has a specific ___________ (or frequency).
The color of the light that is given off depends on how _____ the electron _______ (which depends on how big of a jump it originally made.) The farther the fall, the ___________ energy the photon has.
37. Figure 11.6: Photons of red and blue light.
38. How Light is Produced Since electrons are located only in certain __________ levels (or orbitals) around the nucleus, only certain specific _________ of light are emitted.
Scientists use a _________________ to separate these colors into bands of light. These bands of color look like a ______ code of color which is characteristic of that element. No two elements produce the same ______________ of colors. This can be used to distinguish one element from another contained in a sample. (See Fig. 13.11)
39. Emission Spectrum