AP Chemistry Unit 2 Topics: atomic models: evolutionary evidence, limitations of shell model, evidence of QM idea* Orbital diagrams; how it relates to the periodic table (and properties) e- dot notation isotopes (mass spec): counting p+, no, & e− periodic table: sections, trends using shell model and Coulomb’s Law, PES data *Some of these topics will have a corresponding video from the internet to watch instead of “formal notes”. -- Also, much of Unit 2 contains PowerPoint slides that were taken directly from the 1st year chemistry notes that I use. So if I cover this old material too rapidly, you can listen to the 1st year lecture on my website to refresh your memory of these “review topics”. I will go slower over the “new material” compared to “old material”.

Quantum Mechanical Model http://courses.learn60.ca/mod/book/tool/print/index.php?id=18227 Quantum Mechanical Model

http://www.slideshare.net/laburkett/history-of-the-atom

Quantum Mechanical Model

Bohr’s Energy Levels The energy levels in an atom are sort of like _________ of a ladder. The more energy an electron has, the __________ away from the nucleus it usually will be. The energy levels are not evenly spaced. They get ___________ together as you travel farther away. To move from one “rung” to another requires a “____________” of energy. steps farther closer quantum

Bohr Atomic Model

The difference between continuous and quantized energy levels. continuous energy levels quantized energy levels

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. location address properties Periodic Table how why

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 _______________ ______________ ____________. distance farther principal energy levels 1 2 3 nucleus

shape 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; sublevels spherical ellipsoid

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. d - orbitals criss crossing f - orbitals

How principal energy levels can be divided into sublevels: f s p d s p s

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. orientations 1 3 5 7 spherical 1 ellipsoid 3 4

Quantum Mechanical Model

Principal energy level 2 shown divided into the 2s and 2p sublevels

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. up down arrow spherical up 2

Electron Configurations 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 ____ lowest 2s 3p

Electron Configurations ↑ ↑ Silicon ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ (Energy Level Diagram) ↑ ↓ 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p…

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 ___ 1s 1s Rule #4 (Hund’s Rule): “_______________ rule” or “Bus Seating” rule---> Every “□” in an orbital shape gets an electron before any orientation gets a second e−. Example: □□□ not □□□ 2p 2p 2 ↑ ↓ ↑ ↑ ↓ ↑ opposite ↑ ↓ ↑ ↑ Monopoly ↑ ↑ ↑ ↑ ↓ ↑