Chapter 8 Page 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 = 1s1 or ↑1s

2. Atoms bigger than H are treated by placing additional electrons into H-like orbitals:

Example: He = 1s2 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 = 1s2 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 = 1s22s1 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...

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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 = 1s1 He = 1s2 Li = 1s22s1 Be = 1s22s2

B = 1s22s22p1 ↑↓1s

↑↓2s

↑ 2p

C = 1s22s22p2 ↑↓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 = 1s22s22p3 ↑ ↑ ↑ 2p

O = 1s22s22p4 ↑↓ ↑ ↑ 2p

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F = 1s22s22p5 ↑↓ ↑↓ ↑ 2p

Ne = 1s22s22p6 ↑↓ ↑↓ ↑↓ 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] 3s1 Mg = [Ne] 3s2 Al = [Ne] 3s23p1 Si = [Ne] 3s23p2

P = [Ne] 3s23p3 S = [Ne] 3s23p4 Cl = [Ne] 3s23p5 Ar = [Ne] 3s23p6

14. Note similarity between F and Cl, a look ahead! F = [He] 2s22p5 Cl = [Ne] 3s23p5 same outer electron configuration but different n level 15. Row 4 and 5, the d sublevels start filling: K = [Ar] 4s1 Ca = [Ar] 4s2 Sc = [Ar] 4s23d1 (d starts filling) or [Ar] 3d14s2

Ti = [Ar] 3d24s2 ↑↓4s

↑ ↑ 3d

V = [Ar] 3d34s2 ↑↓4s

↑ ↑ ↑ 3d

1st anomaly: Cr = [Ar] 3d54s1 ↑ 4s ↑ ↑ ↑ ↑ ↑ 3d

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back to normal: Mn = [Ar] 3d54s2 ↑↓ 4s ↑ ↑ ↑ ↑ ↑ 3d

• • •

2nd anomaly: Cu

-using rules one would predict Cu = [Ar] 3d94s2, but in reality, Cu = [Ar] 3d104s1

back to normal: Zn = [Ar] 3d104s2

B. The Periodic Table and Electron Configurations.

1. Electron configuration explains periodicity of element properties.

Examples:

a. Compare alkali metals: Li = [He] 2s1 Na = [Ne] 3s1 K = [Ar] 4s1 Rb = [Kr] 5s1

b. Compare alkaline earths: Mg = [Ne] 3s2 Ca = [Ar] 4s2

c. Compare oxygen and sulfur: O = [He] 2s22p4 S = [Ne] 3s23p4

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)

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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:

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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 = 1s22s22p63s1 Example: Fe = [Ar] 3d64s2 d. Groups in Periodic Table are those having identical outer electron configurations.

3. Noble Gases - chemically inert due to stability of ns2np6 outer configuration. He = 1s2 -filled n = 1 level Ne = 1s22s22p6 -filled n = 2 level Ar = 1s22s22p63s23p6 -ns2np6 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

ns1 ns2 ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 n=1 H n=2 Li Be B C N O F n=3 Na Mg Al Si P S Cl • • • • • • • • • • • • • •

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5. Transition Metals (B Groups): a. All are metals with e- being added to d orbitals.

OUTER CONFIGURATION Sc 4s23dx Zn Y 5s24dx Cd La 6s25dx4f14 Hg

6. Lanthanide and Actinide Series:

a. 4f and 5f are being filled after 1 e- is placed in a d orbital.

6s25d14fx 7s26d15fx

B. Atomic Radii. (Section 8.6)

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.

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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.

I

C. Ionization Energy (IE). (Section 8.6)

1. IE1 = 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. IE2 = amount required to remove second electron.

Ca+1 (g) → Ca+2(g) + e- ↑ add 1145 kJ

IE2 always > IE1 because removing e- from a cation.

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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 8.18.

5. Elements with low IE more likely to form ionic compounds by becoming cations. 6. Monatomic cations > +3 charge are difficult to form: Example: Al3+ stable, Si4+ won’t form.

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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.

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B. Periodicity of the Main Group Elements as exemplified by their Oxides. (Section 8.7) 1. O2 discovered by Priestley (1774).

2 HgO(s) Δ → 2 Hg(l) + O2 (g) mercuric oxide

2. Facts about oxygen:

a. biosphere is 50% oxygen by mass. b. odorless and colorless. c. air = 20% O2, 80% N2, traces of other gases. d. slightly soluble in H2O (enough to sustain marine life).

3. Commercial preparation: distillation of liquid air. 4. O3, ozone, is an unstable allotrope of oxygen:

a. pale blue gas. b. formed by electric spark through O2(g).

5. When O3 decomposes to O2, oxygen atoms O are intermediates. Isolated atoms of

oxygen have unpaired electrons and are called radicals, extremely reactive. 6. Oxygen O2 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. O2 reacts vigorously with Group IA metals to produce:

1.) oxides - Li 2 O−2( )s( )

2.) peroxides - Na2 O−1( )

2 g( ) O22- = peroxides

3.) superoxides - K O−12( )

2 g( ) O2- = superoxide ion

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b. O2 reacts with Group IIA metals at moderate T to form normal oxides

€

M O−2( )

, and

at high O2 pressure to form peroxides

€

M O−1( )

2 with the heavier IIA metals. c. O2 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.

Na2O(s) + H2O → 2 NaOH metal oxide metal hydroxide (a basic anhydride) e. O2 reacts with nonmetals to form molecular compounds.

€

4C(0)(s)+O2(g)→ C

(+4)O2(g)

f. These nonmetal oxides are acidic oxides and react with water to form acids. Thus

called acid anhydrides. CO2(g) + H2O → H2CO3(aq) nonmetal oxide ternary acid (acid anhydride)

SO2(g) + H2O → H2SO3(aq) (no change in SO3(g) + H2O → H2SO4(aq) ox state of N2O5(s) + H2O → 2 HNO3(aq) nonmetal) P4O10(s) + 6 H2O → 4 H3PO4(aq)

7. Reactions of metal oxides with nonmetal oxides form salts.

CaO(s) + SO3(g)→ CaSO4(s) 6 Na2O + P4O10 → 4 Na3PO4(s)

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8. Combustion reactions = redox reaction in which oxygen combines rapidly with any

oxidizable materials.

CH4 + 2 O2 → CO2 + 2 H2O + heat 9. Air pollution.

a. Combustion of any materials containing sulfur produces SO2(g). b. Slowly converts to SO3(g) in atmosphere. 2 SO2(g) + O2(g) → 2 SO3(l) c. Combines with H2O to produce acid rain. SO3 + H2O → H2SO4 d. Combustion of any materials containing nitrogen produces NO, nitric oxide. N2 + O2 → 2 NO(g) 2 NO(g) + O2(g)

uv light → 2 NO2(g) brown gas 3 NO2(g) + H2O → 2 HNO3(aq) + NO acid rain

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Notes: