Quantum numbers are used to describe where an e – is in an atom.

1. Quantum numbers are used to describe where an e– is in an atom. (n = integers 1, 2, 3,...) 1. principal quantum number, n -- correspond to the energy level of the electrons -- All orbitals having the same n are called an electron shell (e.g., 2s and 2p). 2. angular momentum quantum number, l (l = integers from 0 up to (n – 1)) -- This number defines the type of subshell: s = 0, p = 1, d = 2, f = 3 -- For a given shell, the energies of orbitals go: s < p < d < f

2. Quantum numbers (cont.) 3. magnetic quantum number, ml (ml = integers from –l to l) -- describes the orientation of an orbital in space -- You should know the shapes and orientations of the s, p, and d orbitals: s, px, py, pz, dyz, dxz, dxy, dx2– y2, dz2 shapes and orientations of d orbitals

3. Wolfgang Pauli (1900–1958) 4. electron spin quantum number, ms -- only two values: +½ or –½ (“spin-up” and “spin-down”) -- Pauli exclusion principle: No two electrons in an atom may have the same set of four quantum numbers (i.e., an orbital may hold only two electrons, and they must have opposite spins).

4. (a) Predict the number of subshells in the third shell. n = 3, so… l could be 0, 1, or 2 3 subshells (b) Give the label for each of these subshells. For: l = 0… 3s l = 1… 3p l = 2… 3d (c) How many orbitals are in each of these subshells? For: l = 0… ml = 0 3s: 1 orbital ml = –1, 0, +1 l = 1… 3p: 3 orbitals ml = –2, –1, 0, +1, +2 l = 2… 3d: 5 orbitals

5. What are the values of n and l for the following sublevels? 2s 3d 4p 5s 4f n = 2 n = 3 n = 4 n = 5 n = 4 l = 0 l = 2 l = 1 l = 0 l = 3 Write the possible sets of the four quantum numbers for a 4p electron. 4p orbitals ml : –1 0 +1 4, 1, –1, +½ 4, 1, 0, +½ 4, 1, +1, +½ 4, 1, –1, –½ 4, 1, 0, –½ 4, 1, +1, –½

6. Write the four quantum numbers of each of the six 3d electrons of an iron atom. 3d orbitals ml : –2 –1 0 +1 +2 3, 2, –2, +½ ? 3, 2, –1, +½ 3, 2, 0, +½ 3, 2, +1, +½ 3, 2, +2, +½ 3, 2, ?, –½

7. lustrous (shiny) malleable (can hammer into shape) ductile (can pull into wire) good conductors (heat and electricity) Brief Review of the Periodic Table metals: left side of Table; form cations properties:

8. Brief Review of the Periodic Table (cont.) nonmetals: right side of Table; form anions good insulators gases or brittle solids properties: neon sulfur iodine bromine Ne S8 I2 Br2

9. nonmetals metals computer chips Si and Ge computer chips Ge and Si  Brief Review of the Periodic Table (cont.) metalloids (semimetals): “stair” between metals and nonmetals (B, Si, Ge, As, Sb, Te, Po, At) properties: in-between those of metals and nonmetals; “semiconductors”

10. group 1 (except H); 1+ charge; very reactive alkali metals: group 2; 2+ charge; less reactive than alkalis alkaline earth metals: groups 3–12; variable charges transition elements: group 16; 2– charge; reactive chalcogens: group 17; 1– charge; very reactive halogens: noble gases: group 18; no charge; unreactive lanthanides: elements 58–71 contain f orbitals actinides: elements 90–103 groups 1, 2, 13–18 main block (representative) elements:

11. What family of elements has an ns2 valence electron configuration? alkaline earth metals

12. “RuRh…!!” Anomalies in the Electron Configurations Your best guide to writing e– configs is “The Table,” but there are a few exceptions. e.g., Cr: [ Ar ] 4s1 3d5 Cu: [ Ar ] 4s1 3d10 These exceptions are due to the closeness in energy of the upper-level orbitals. Other exceptions are… Mo, Ru, Rh, and Ag. All of these exceptions have a single valence-level s electron.