Part One: Electronic Structure Of Atoms. 1. Electron configuration is shorthand notation for what AO the electron occupies:

Transcription

1 CHAPTER EIGHT: ELECTRON CONFIGURATIONS AND PERIODICITY Part One: Electronic Structure Of Atoms A. Electron Configurations of Multi-Electron Atoms. (Section 8.1) 1. Electron configuration is shorthand notation for what AO the electron occupies: Example - The ground state of H atom (lowest energy state): H = 1s 1 or 1s 2. Atoms bigger than H are treated by placing additional electrons into H-like orbitals: Example: He = 1s 2 or 1s 3. Note that maximum of 2 electrons can go into an atomic orbital with opposite spins, consequence of Pauli Principle. 4. Pauli Exclusion Principle = no two e - in an atom can have identical set of 4 quantum numbers. Example: He = 1s 2 or 1s has two electrons in states: (1, 0, 0, +1/2) (1, 0, 0, -1/2) 5. P.E.P. also implies that for 3 e - atom like Li, would have to start to fill 2s AO: Li = 1s 2 2s 1 or 1s 2s 6. Aufbau Principle (Section 8.2) = Building-up principle = electron configuration of multi e - atoms built up by addition of electrons to H-like AO to give the lowest total energy for the atom. filling order = 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p... Chapter 8 Page 1

2 7. In multi-electron atoms the energies of AO s follow this design: 8. Note that 4s is lower in energy than 3d, so 4s fills first. 9. Now let s write some electron configurations. H = 1s 1 He = 1s 2 Li = 1s 2 2s 1 Be = 1s 2 2s 2 B = 1s 2 2s 2 2p 1 1s 2s 2p C = 1s 2 2s 2 2p 2 1s 2s 2p 10. Hund s Rule = when filling a sublevel having more than one AO (such as 2p sublevel) one places electrons singly in separate orbitals before pairing begins. These unpaired e - have parallel spins. This is lower energy. 11. Let s continue: N = 1s 2 2s 2 2p 3 2p O = 1s 2 2s 2 2p 4 2p Chapter 8 Page 2

3 F = 1s 2 2s 2 2p 5 2p Ne = 1s 2 2s 2 2p 6 2p 12. Note at Neon, 2p sublevel is filled, and also n = 2 level is filled. At He n = 1 level is filled. These are stable, chemically inert configurations. 13. Let s do Row 3: Na = [Ne] 3s 1 Mg = [Ne] 3s 2 Al = [Ne] 3s 2 3p 1 Si = [Ne] 3s 2 3p 2 P = [Ne] 3s 2 3p 3 S = [Ne] 3s 2 3p 4 Cl = [Ne] 3s 2 3p 5 Ar = [Ne] 3s 2 3p Note similarity between F and Cl, a look ahead! F = [He] 2s 2 2p 5 Cl = [Ne] 3s 2 3p 5 same outer electron configuration but different n level 15. Row 4 and 5, the d sublevels start filling: K = [Ar] 4s 1 Ca = [Ar] 4s 2 Sc = [Ar] 4s 2 3d 1 (d starts filling) or [Ar] 3d 1 4s 2 Ti = [Ar] 3d 2 4s 2 V = [Ar] 3d 3 4s 2 4s 4s 3d 3d 1st anomaly: Cr = [Ar] 3d 5 4s 1 4s 3d Chapter 8 Page 3

4 back to normal: Mn = [Ar] 3d 5 4s 2 4s 3d 2nd anomaly: Cu -using rules one would predict Cu = [Ar] 3d 9 4s 2, but in reality, Cu = [Ar] 3d 10 4s 1 back to normal: Zn = [Ar] 3d 10 4s 2 B. The Periodic Table and Electron Configurations. 1. Electron configuration explains periodicity of element properties. Examples: a. Compare alkali metals: Li = [He] 2s 1 Na = [Ne] 3s 1 K = [Ar] 4s 1 Rb = [Kr] 5s 1 b. Compare alkaline earths: Mg = [Ne] 3s 2 Ca = [Ar] 4s 2 c. Compare oxygen and sulfur: O = [He] 2s 2 2p 4 S = [Ne] 3s 2 3p 4 2. Therefore, electron configurations explain the shape of the Periodic Table, and vice versa, we can use Periodic Table to figure out electron configuration. (see Figure 8.12 in text) Chapter 8 Page 4

5 C. Magnetic properties of atoms. 1. Electron in an atom behaves like a tiny magnet and orients in a magnetic field. 2. Magnetic field from two electrons with paired spins (one up and one down) cancels itself out. No net magnetism. 3. Atoms with all paired-up electrons are called diamagnetic - they are not attracted by an external magnetic field, but actually slightly repelled. Example: Hg vapor Why? 4. Unpaired electrons in an atom impart an overall magnetism to the atom and these are called paramagnetic. They are attracted by a magnetic field. Example: Na vapor Why? 5. This behavior proves Hund s Rule is in effect. For example, the electronic configuration of Carbon is: Chapter 8 Page 5

6 Part II: Periodic Properties of the Elements A. Theoretical Foundation of the Periodic Law. (Section 6.1) 1. Periodic law = properties of the elements are periodic (repeating) functions of their atomic numbers. 2. Theoretical basis: a. Outer electrons of an atom largely determine its properties. b. Outer electrons are called valence electrons. c. Outer electrons are those with the highest n quantum number. Example: Na = 1s 2 2s 2 2p 6 3s 1 Example: Fe = [Ar] 3d 6 4s 2 d. Groups in Periodic Table are those having identical outer electron configurations. 3. Noble Gases - chemically inert due to stability of ns 2 np 6 outer configuration. He = 1s 2 Ne = 1s 2 2s 2 2p 6 Ar = 1s 2 2s 2 2p 6 3s 2 3p 6 -filled n = 1 level -filled n = 2 level -ns 2 np 6 configuration is extremely stable and chemically inert. 4. Representative Block Elements (A Groups): a. Have partially occupied outer level. b. Last electron to be added in Aufbau procedure was added to an s or p orbital. OUTER LEVEL CONFIGURATION IA IIA IIIA IVA VA VIA VIIA ns 1 ns 2 ns 2 np 1 ns 2 np 2 ns 2 np 3 ns 2 np 4 ns 2 np 5 n=1 H n=2 Li Be B C N O F n=3 Na Mg Al Si P S Cl Chapter 8 Page 6

7 5. Transition Metals (B Groups): a. All are metals with e - being added to d orbitals. OUTER CONFIGURATION Sc 4s 2 3d x Zn Y 5s 2 4d x Cd La 6s 2 5d x 4f 14 Hg 6. Lanthanide and Actinide Series: a. 4f and 5f are being filled after 1 e - is placed in a d orbital. 6s 2 5d 1 4f x B. Atomic Radii. (Section 8.6) 7s 2 6d 1 5f x 1. Size of atoms determined by size of electron cloud around the nucleus. This size is somewhat indefinite. 2. Size dictates how densely atoms pack w/ other atoms in solids. Chapter 8 Page 7

8 3. Within a Group A series, atomic radii increase from top to bottom of periodic chart: small Li < Na < K < Rb < Cs large 4. Moving across a period, atomic radii decrease. C. Ionization Energy (IE). (Section 8.6) 1. IE 1 = 1st ionization energy = minimum amount of energy required to remove most loosely held electron from gaseous neutral atom. Ca(g) Ca +1 (g) + e - add 590 kj 2. IE 2 = amount required to remove second electron. Ca +1 (g) Ca +2 (g) + e - add 1145 kj IE 2 always > IE 1 because removing e - from a cation. I Chapter 8 Page 8

9 3. Trend: (goes opposite Atomic Radii) 4. Measures how tightly bound the outermost electrons are. a. When IE low, e - easy to remove. b. Metals have low IE, nonmetals high IE. c. See Figure Elements with low IE more likely to form ionic compounds by becoming cations. 6. Monatomic cations > +3 charge are difficult to form: Example: Al 3+ stable, Si 4+ won t form. Chapter 8 Page 9

10 7. Successive IE s. D. Electron Affinity (EA). (Section 6-4) 1. EA = amount of energy needed to attach an electron to gaseous neutral atom. 2. Cl(g) + e - Cl - (g) EA = -348 kj 348 kj released Be(g) + e - Be - (g) EA 0 0 kj absorbed 3. Elements with very negative EA gain e - easily to form anions. (i.e. nonmetals) 4. See Table 8.4. Chapter 8 Page 10

11 B. Periodicity of the Main Group Elements as exemplified by their Oxides. (Section 8.7) 1. O 2 discovered by Priestley (1774). 2. Facts about oxygen: 2 HgO(s) mercuric oxide Δ 2 Hg(l) + O 2 (g) a. biosphere is 50% oxygen by mass. b. odorless and colorless. c. air = 20% O 2, 80% N 2, traces of other gases. d. slightly soluble in H 2 O (enough to sustain marine life). 3. Commercial preparation: distillation of liquid air. 4. O 3, ozone, is an unstable allotrope of oxygen: a. pale blue gas. b. formed by electric spark through O 2 (g). 5. When O 3 decomposes to O 2, oxygen atoms O are intermediates. Isolated atoms of oxygen have unpaired electrons and are called radicals, extremely reactive. 6. Oxygen O 2 combines readily with all other elements to form oxides except noble gases (Group VIIIA elements) and noble metals (Au, Pd, Pt). 7. Oxygen combines with metals in general to form basic oxides or amphoteric oxides. (one that has both acidic and basic properties) a. O 2 reacts vigorously with Group IA metals to produce: ( ) s ( ) 2 g 1.) oxides - Li 2 O 2 ( ) 2.) peroxides - Na 2 O 1 ( ) 3.) superoxides - K O 1 2 ( ) O 2 2- = peroxides - 2( g) O 2 = superoxide ion Chapter 8 Page 11

12 b. O 2 reacts with Group IIA metals at moderate T to form normal oxides M ( O 2), and at high O 2 pressure to form peroxides M ( O 1) 2 with the heavier IIA metals. c. O 2 reacts with all other metals (except noble) to form the oxides with the normal stoichiometry. d. These metal oxides react readily with water to form bases. Thus called basic anhydrides. Na 2 O(s) + H 2 O 2 NaOH metal oxide metal hydroxide (a basic anhydride) e. O 2 reacts with nonmetals to form molecular compounds. 4 (0) C (s)+ O 2 (g) (+4) C O 2 (g) f. These nonmetal oxides are acidic oxides and react with water to form acids. Thus called acid anhydrides. CO 2 (g) + H 2 O H 2 CO 3 (aq) nonmetal oxide ternary acid (acid anhydride) SO 2 (g) + H 2 O H 2 SO 3 (aq) SO 3 (g) + H 2 O H 2 SO 4 (aq) N 2 O 5 (s) + H 2 O 2 HNO 3 (aq) P 4 O 10 (s) + 6 H 2 O 4 H 3 PO 4 (aq) 7. Reactions of metal oxides with nonmetal oxides form salts. CaO(s) + SO 3 (g) CaSO 4 (s) 6 Na 2 O + P 4 O 10 4 Na 3 PO 4 (s) (no change in ox state of nonmetal) Chapter 8 Page 12

13 8. Combustion reactions = redox reaction in which oxygen combines rapidly with any oxidizable materials. CH O 2 CO H 2 O + heat 9. Air pollution. a. Combustion of any materials containing sulfur produces SO 2 (g). b. Slowly converts to SO 3 (g) in atmosphere. 2 SO 2 (g) + O 2 (g) 2 SO 3 (l) c. Combines with H 2 O to produce acid rain. SO 3 + H 2 O H 2 SO 4 d. Combustion of any materials containing nitrogen produces NO, nitric oxide. N 2 + O 2 2 NO(g) 2 NO(g) + O 2 (g) uv light 2 NO 2 (g) brown gas 3 NO 2 (g) + H 2 O 2 HNO 3 (aq) + NO acid rain Chapter 8 Page 13

14 Notes: Chapter 8 Page 14