Electrons! Chapter 5, Part 2
3. Contained within sublevels are orbitals: pairs of electrons each having a different space or region they occupy a. Each sublevel contains certain orbitals: i. s sublevel contains 1 s orbital ii. p sublevel contains 3 p orbitals (p x, p y, p z ) iii.d sublevel contains 5 d orbitals (d xy, d xz, d yz, d x 2-y2, d z 2) iv.f sublevel contains 7 f orbitals v. Each orbital holds 2 electrons b. Orbital designations follow this pattern: 1s, 2s, 2p, 3d, etc. The number is the principal energy level, the letter is the type of orbital 4. Only specific sublevels & orbitals are contained in each principal energy level a. Principal energy level 1 contains 1s b. Principal energy level 2 contains 2s, 2p c. Principal energy level 3 contains 3s, 3p, 3d d. Principal energy level 4 contains 4s, 4p, 4d, 4f e. Principal energy levels 5+ contain s, p, d, f
Copy Table 5-2 (p. 134)
III. Electron Configurations A.Ground-State Configurations 1. Electron configuration: arrangement of electrons in an atom 2. There are 3 rules that dictate how electrons are arranged a. aufbau principle: each electron occupies the lowest energy orbital available i. aufbau diagram shows the sequence of orbitals from lowest to highest energy; key points: All orbitals related to an energy sublevel have equal energy In a multi-electron atom, the energy sublevels within a principal energy level have different energies In order of increasing energy, the sequence of energy sublevels within a principal energy level is s, p, d, f Orbitals related to energy levels within one principal energy level can overlap orbitals related to energy sublevels within another principal energy level
Copy Figure 5-17 (p. 135)
b. Pauli exclusion principle: 2 electrons may occupy a single orbital, but only the electrons have opposite spins i. An electron can only spin one way -- clockwise ( ) or counterclockwise ( ) ii. An orbital must have one clockwise & one counter-clockwise electron c. Hund s rule: electrons in the same sublevel do not pair until each orbital is occupied with one electron (sharing rule) i. In other words, each orbital must have one electron with a clockwise spin before any electrons with counterclockwise spins will be added
Homework: p. 134 #13-17
B. Representing Electron Configuration 1. Orbital Diagrams a. One box per orbital -- each box is labeled with the name of the orbital (1s, 2s, etc.) b. Electrons are represented by arrows (# e - = atomic #) c. Key: empty orbital orbital with one electron d. Example: oxygen filled orbital with 2 electrons
2. Electron configuration notation a. Includes the orbital designation (1s, 2s, 2p, etc.) with a superscript indicating the number of electrons in that orbital b. Examples: hydrogen 1s 1 helium 1s 2 oxygen 1s 2 2s 2 2p 4 neon 1s 2 2s 2 2p 6 sodium 1s 2 2s 2 2p 6 3s 1 c. To represent longer electron configurations, noble-gas configurations can be used i. Notation of noble gas in previous period and the notation of the energy level being filled ii. Example: sodium [Ne]3s 1 magnesium [Ne]3s 2
Copy: Table 5-3 (p. 137) Table 5-4 (p. 138)
3. Diagonal Rule a. Easy way to determine the correct electron configuration & the order in which orbitals fill up **NOTE: there are some exceptions to the diagonal rule (Cu, Cr) -- see PT
C. Valence Electrons 1. Definition: electrons in an atom s outermost orbitals; usually correspond to electrons in highest energy level 2. Valence electrons determine if & how an atom will react 3. General rule: eight valence electrons make an atom chemically stable 4. Electron dot structures show the number & arrangement of valence electrons
5. How to draw electron dot structures: a. Determine how many valence electrons are in the element -- do this by writing out the element s noble-gas configuration i. Example: sulfur [Ne]3s 2 3p 4 b. Write the element s symbol -- this represents the name of the element, the nucleus, & all of its inner-level electrons c. Place one dot at a time on each side of the element symbol, then pair up until total number is reached i. Example: sulfur
Copy: Table 5-5 (p. 140) Practice Problems: p. 139 #18-22 Homework: p. 141 #24-28