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1 Chapter 17 Many-Electron Atoms and Chemical Bonding 17.1 Many-Electron Atoms and the Periodic Table 17.2 Experimental Measures of Orbital Energies 17.3 Sizes of Atoms and Ions 17.4 Properties of the Chemical Bond 17.5 Ionic and Covalent Bonds 17.6 Oxidation States
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Page 1: Chapter 17 Many-Electron Atoms and Chemical Bonding · Chapter 17 Many-Electron Atoms and Chemical Bonding 17.1 Many-Electron Atoms and the Periodic Table ... 19 Filling the orbitals

1

Chapter 17 Many-Electron Atoms andChemical Bonding

17.1 Many-Electron Atoms and the Periodic Table

17.2 Experimental Measures of Orbital Energies

17.3 Sizes of Atoms and Ions

17.4 Properties of the Chemical Bond

17.5 Ionic and Covalent Bonds

17.6 Oxidation States

Page 2: Chapter 17 Many-Electron Atoms and Chemical Bonding · Chapter 17 Many-Electron Atoms and Chemical Bonding 17.1 Many-Electron Atoms and the Periodic Table ... 19 Filling the orbitals

2

17.1 Many-Electron Atoms

Many electron atoms and the periodic table

Building up electron configurations

Building up from H to Ar: Pauli and Aufbau principles

Building from K to Kr: Transition elements and d orbitals

Electron shells, effective nuclear charge

Hund’s rule: paramagnetism and diamagnetism.

Correlations of periodic properties of familes ofelements and electron configurations

Page 3: Chapter 17 Many-Electron Atoms and Chemical Bonding · Chapter 17 Many-Electron Atoms and Chemical Bonding 17.1 Many-Electron Atoms and the Periodic Table ... 19 Filling the orbitals

3

Multielectron atoms:

Every electron in a multielectron atom is assignedfour quantum numbers (n, l, ml and ms) that uniquely

define its chemical and physical properties.

From the use of quantum numbers, we can envisionevery electron of a multielectron atom in terms of a

characteristic energy (En), size (r), shape (l),orientation (ml) and spin (ms).

The highest energy (valence) electrons are ofgreatest chemical interest

Page 4: Chapter 17 Many-Electron Atoms and Chemical Bonding · Chapter 17 Many-Electron Atoms and Chemical Bonding 17.1 Many-Electron Atoms and the Periodic Table ... 19 Filling the orbitals

4

u

s

6.941

F7 8 9 105

14.0067 20.179

Ne18.9984

O

He2

ArAl14 15 16 17 18

P13

30.9738 35.453

S

KrGa32 33 34 35 36

As31

74.9216 79.904

Se

XeIn50 51 52 53 54

Sb49

121.757 126.905

Te

RnTl82 83 84 85 86

Bi81

208.980 (210)

Po

LuDy67 68 69 70 71

Er66

167.26 173.04

Tm

ZnMn26 27 28 29 30

Co25

58.9332 63.546

Ni

CdTc44 45 46 47 48

Rh43

102.906 107.868

Pd

HgRe76 77 78 79 80

Ir75

192.22 196.967

Pt

MtRf105 106 [107] [108] [109]

Sg104

(263) (265)Ns

CrK20 21 22 23 24

Sc19

44.9559 50.9415

Ti

MoRb38 39 40 41 42

Y37

88.9059 92.9064

Zr

WCs56 57 72 73 74

La55

138.906 180.948

Hf

EuCe59 60 61 62 63

Nd58

144.24 150.36

Pm

AmTh91 92 93 94 95

U90

238.029 (244) (252)

NoFmEsCf99 100 10198

(257) (260)(259)

Gd Tb64 65

Cm Bk96 97

(247)

Li3 4

15.9994

C NB6

10.81

Cl28.0855

4.00260

Si32.066 39.948

Br72.59

26.9815

Ge78.96 83.80

I118.710

69.72

Sn127.60 131.29

At207.2

114.82

Pb(209) (222)

Yb164.930

204.383

Ho168.934 174.967

Cu55.847

162.50

Fe58.69 65.39

Ag101.07

54.9380

Ru106.42 112.41

Au190.2

(98)

Os195.08 200.59

Hs(262)

186.207

Ha(262) (268)

V40.078

(261)

Ca47.88 51.996

Nb87.62

39.0983

Sr91.224 95.94

Ta137.33

85.4678

Ba178.49 183.85

Sm140.908

132.905

Pr(145) 151.96

Pu231.036

140.12

Pa237.048 (243)

Np232.038 (258)

Lr102 103

Md(251)

158.925157.25

(247)

Be9.01218

Na11

Mg12

Fr87

(223)

H1

1.0079

12.011

89

s Actinide Series

u Lanthanide Series

88

AcRa227.028226.025

The Periodic Table of the Elements

24.30522.9898

1a

2a

3a 4a 5a 6a 7a 8 8 8 1b 2b

3a 4a 5a 6a 7a

8a

Building up the table from electronconfigurations: First four periods (H-Kr)

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Atoms of the various elements differ from each other intheir values of Z (atomic number) and electrons.

Electronic structure of atoms of the elements:

Electrons in atoms are arranged in orbitals, shells andsubshells. Orbitals having the same value of n aresaid to be in the same shell. Orbitals having thesame values of n and l are said to be in the samesubshell.

Orbitals of all elements are hydrogen like and arecharacterized by the quantum numbers n, l and ml.

Energies of orbitals depend on (n + l), not just n.

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The orbital approximation: The electron density ofan isolated many-electron atom is approximately the

sum of the electron densities of each of theindividual electrons taken separately.

For atoms with more than one electron,approximations are required in order to make

quantitative quantum mechanical approximations.

The approximation amounts to treat each electron as if itwere moving in a field o charge that is the net result ofthe nuclear attraction and the average repulsions of all

the other electrons.

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Building up the ground state electron configurationsof atoms

An electron configuration is a list of the occupiedorbitals and the number of electrons in each.

The electron configuration of lowest energy istermed the ground state electronic configuration.

The ground state electron configuration is built by fillingthe lowest energy orbitals first (Aufbau principle) and

obeying the Pauli principle and Hund’s rule

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Orbitals: the individual components of a shell or subshellExample: the px, py and pz orbitals are the components of any

p (l = 1) subshell and each orbital has the same energy

Shells, subshells and orbitals: definitions

Shell: a collection of orbitals with the same value of n

Example: the three orbitals 3s, 3p, 3d comprise a shell with n = 3

Subshell: a collection of orbitals with the same value of n and l.The orbital of a subshell have the same energy for the same

value of l.

Example: for the n = 3 shell there are three subshells, the 3ssubshell (l = 0), the 3p subshell (l = 1), and the 3d subshell (l = 2),

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En = -(Zeff2/n2)Ry = energy of electron in orbit

rn = (n2/Zeff)a0 = radius of a Bohr orbitReplace Z (actual charge) with Zeff (effective charge)

The Bohr one electron atom as a starting point for theelectron configurations of multielectron atoms.

Structure of multielectron atoms:

Quantum numbers of electrons: n, l, ml, msElectron configurations: 1sx2sx2px3sx3px, etc (x = number of electrons)Core electrons and valence electrons (Highest value of n)

Some periodic properties of atoms we shall study:

Energy required to remove and add an electron (En)Size of atoms (rn)

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The energy of an orbital of a hydrogen atom or any one electronatom only depends on the value of n

shell = all orbitals with the same value of nsubshell = all orbitals with the same value of n and l

an orbital is fully defined by three quantum numbers, n, l, and ml

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For a multielectron atom, the energy of subshellsincrease with l for a given value of n

Example: f orbitals

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Determining the ground state electronic configuration ofneutral atoms

(1) Use the (n + l) rule to determine the relativeenergies of the atomic orbitals from 1s to nl

(2) Imagine a bare nucleus of charge +Z surroundedby empty atomic orbitals.

(3) Add Z electrons to the empty orbitals starting withthe lowest energy orbital first, obeying the Pauli principleas required.

(4) Electrons are placed in orbitals of lowest energy according tothe Aufbau principle to obtain the ground state electronconfiguration. Apply other rules as required.

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Orbital (n + l) Comment

1s (1 + 0 = 1)2s (2 + 0 = 2)2p (2 + 1 = 3)3s (3 + 0 = 3) Lower n (2p versus 3s) has lower energy3p (3 + 1 = 4)4s (4 + 0 = 4) Lower n (3p versus 4s) has lower energy3d (3 + 2 = 5)4p (4 + 1 = 5) Lower n (3d versus 4p) has lower energy5s (5 + 1 = 6)

The (n + l) rule: The ordering of the energies of the orbitals in amultielectron atom increases with the value of n + l.

Sub-rule: When two orbitals of different n have the same value of n + lthe orbital with the lower value of n has the lower energy state.

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Relative orbital energies for the multielectron atom.The energy of an orbital of a multielectron atom

depends on n and l (but not ml)

2s < 2p < 3s

3s < 3p < 3d ~ 4s (may switch with Z)

Note energy levels are getting closertogether for n = 3 as expected fromthe Bohr atom.This means that factors ignored mayhave to be considered

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Atomic Energy Levelsaccording to the

(n + l) rule

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Ground state electron configuration of a many electronatom: Governs reactivity of atoms under normal

condition

(1) Imagine a bare nucleus of charge +ZImagine empty orbitals surrounding the nucleus

(2) Fill the orbitals with Z electrons for the neutralatom following two principles:

(3) Aufbau principle: fill lowest energy orbitals first

(4) Pauli exclusion principle: each electron must havefour different quantum numbers (maximum of 2

electrons in an orbital).

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The Aufbau principle.

Systems in nature tend to minimize their energy.

The ground state of an atom is the one with the lowest energy.

All other allowed states are excited states of the atom.

The Aufbau principle for atoms: The systematic adding electronsto orbitals of an atom (ZA) with atomic number Z, beginning withthe orbital of lowest energy and continuing to add electrons, in

accordance with the Pauli principle, until Z electrons are added tothe atom.

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Constructing the periodic table by filling orbitals with electrons(electron configurations) in accordance with the Pauli and Aufbau

principles.

Aufbau: Fill 1s orbital firstPauli: no more than two electronsin the 1s orbitalThe basis of the duet rule:filling a shell 1s subshell filledwith 2He = stable electron coregiven symbol [He].

Construction of the first row of the periodic table.Electron configurations: 1H and 2He.

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Filling the orbitals of 3Li, 4Be and 5B

Aufbau: Fill 1s orbitalfirst, then 2s, then 2p.Pauli: no more than two

electrons in the 1s orbital.

2s subshell filled with 4Be.

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For carbon, how do the two 2pelectrons distributethemselves in the three 2porbitals?

For nitrogen, how do the three2 p electrons distributethemselves in the three 2porbitals?

Filling the orbitals of 6C and 7N. The need for a third rule(Hund’s rule):

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Hund’s rule: Applies when filling the orbitals of a subshellwith electrons (np or nd or nf subshells). Or moregenerally when filling orbitals of identical energy

When adding electrons to a subshell, the ground stateelectronic configuration is formed by maximizing thenumber of electrons with parallel spins (↑)(↑) before

pairing two electrons in one orbital (↑↓)().

Example: 6C = [He]2s22px(↑)2py(↑)2pz() = ground state

Example: 6C = [He]2s22px(↑↓)2py()2pz() = excited state

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Filling the orbitals of 6C and 7N. The need for a third rule(Hund’s rule):

Hund’s Rule: When electrons occupyorbitals of the same energy, the

lowest energy state corresponds tothe configuration with the greatest

number of orbitally and spinunpaired electrons.

When the configuration iswritten as 1s22s22p2 it is

understood that twodifferent p orbitals (px, py)

are occupied.

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The presence of two orbitally and spin unpairedelectrons in the ground state of carbon makes the

atom paramagnetic.

A paramagnetic substance is attracted to a magneticfield. A diamagnetic substance is repelled from a

magnetic field.

All substances which possess one or more orbitallyunpaired electrons are paramagnetic.

All substances which possess only spin pairedelectrons are diamagnetic.

Paramagnetic and diamagnetic substances

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Examples of diamagnetic and paramagnetic atoms

Which of the following atoms are paramagnetic?

1H, 2He, 3Li, 4Be, 5B, 6C, 7N, 8O, 9F, 10Ne

1H, 3L, 5B, 7N, 9F must be paramagnetic since theypossess an odd number of electrons.

4Be, 6C and 8O paramagnetic or diamagnetic?

2He and 10Ne are diamagnetic (filled shells).

4Be: 1s22s2 Two filled shells: diamagnetic6C: 1s22s2p2 Hund’s rule: px(↑)py(↑)8O: 1s22s2p4 Let’s see

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Are the following two configurations allowed for 7Nby the Pauli principle?

1s22s22px12py

12pz1 or 1s22s22px

22py12pz

0

The Pauli exclusions principle does not forbid the existence of anyof the Pauli allowed configurations.

Ans: Yes. Which is more stable?

1s22s22px(↑)2py(↑)2pz(↑) is more stable than1s22s22px(↑↓)2py(↑)2pz() by Hund’s rule

If there are more than electron configurations one allowed Pauliconfiguration, the lowest energy one (ground state) will bepredicted by Hund’s rule and the others will be excited states.

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Filling the orbitals of 8O, 9F and 10Ne

Filling the 2p subshellproduces another stableconfiguration ofelectrons which servesas the core shell of thethird row: symbol [Ne]

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Atom Configuration Magnetic PropertiesCore/Valence electrons

3Li [Ne]2s (↑) Paramagnetic4Be [Ne] 2s2 (↑↓) Diamagnetic (Closed shell)5B [Ne] 2s22p1 (↑) Paramagnetic6C [Ne] 2s22p2 (↑↑) Paramagnetic7N [Ne] 2s22p3 (↑↑↑) Paramagnetic8O [Ne] 2s22p4 (↑↓↑↑) Paramagnetic9F [Ne] 2s22p5 (↑↓↑↓↑) Paramagnetic10Ne [Ne] 2s22p6 (↑↓↑↓↑↓) Diamagnetic (Closed shell)

How do electronic configurations connect with valence electrons andLewis structures?

Building up the Periodic TableFrom 3Li to 10Ne: Paramagnetism and Diamagnetism

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Correlation of valence electron and Lewis structures: 2pindicates an unpaired electron in a 2p orbital

N 2[He]2s22px2py2pz

O 2[He]2s22px22py2pz

F 2[He]2s22px22py

22pz

Ne 2[He]2s22px22py

22pz2

Filled shell

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Building up the third row of the periodic table:From Na to Ar: Paramagnetism and Diamagnetism

Atom Configuration Magnetic properties

Core/Valence electrons11Na [Ne] 3s (↑) Paramagnetic12Mg [Ne] 3s2 (↑↓) Diamagnetic (Closed shell)13Al [Ne] 3s23p1 (↑) Paramagnetic14Si [Ne] 3s23p2 (↑↑) Paramagnetic15P [Ne] 3s23p3 (↑↑↑) Paramagnetic16S [Ne] 3s23p4 (↑↓↑↑) Paramagnetic17Cl [Ne] 3s23p5 (↑↓↑↓↑) Paramagnetic18Ar [Ne] 3s23p6 (↑↓↑↓↑↓) Diamagnetic (Closed shell)

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So far, so good, but now for something totally different:The transition metals: Sc through Zn. Filling the d orbitals

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Electron configuration of the transitionelements:

21Sc through 30Zn

The d block elements:21Sc, 22Ti, 23V, 24Cr, 25Mn, 26Fe, 27Co, 28Ni, 29Cu, 30Zn

d orbitals raise their ugly lobes!

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Atom Configuration

19K (Group I) 18[Ar]4s20Ca (Group II) 18[Ar]4s2

_________________________________ d orbitals fill up31Ga (Group III) 18[Ar] 4s23d104p1

32Ge (Group IV) 18[Ar] 4s23d104p2

33As (Group V) 18[Ar] 4s23d104p3

34Se (Group VI) 18[Ar] 4s23d104p4

35Cl (Group VII) 18[Ar] 4s23d104p5 36Kr (Group VIII) 18[Ar] 4s23d104p6

The fourth row of the periodic table: From 19K to 36KrFrom the results of the second and third row we expect for the

representative elements of Groups I-VIII:

How do the d orbital fill up for 21Sc through 30Zn?

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Building the third full row of the periodic table:19K-36Kr

The 4s orbital is slightlymore stable than the 3dorbital at the beginning ofthe third full period of theperiodic table:

19K = [Ar]4s13d0

20Ca = [Ar]4s23d0

The reason is that the 4s orbital has a higher probability ofbeing closer to the nucleus and see a greater effective Zeff

than a 3d orbital.

The 4s and 3d orbitals are close in energy in the one electron atom.Difficult to predict stability for multielectron atom.

Let’s take a look at how the core electrons screen Z.

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Effectivenuclear charge(Eeff) for 12Mg =is Z(Mg) - [Ne]core electrons =12 = 10 = 2+

Mg = 10[Ne]3s2

Effective nuclear charge (Eeff) is the nuclear charge experiencedby an outer (valence) electron after the shielding due to theshielding of the inner (core) electrons are taken into account

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Electron shielding (screening) of the nuclear charge byother electrons

Why is the energy of a 3s orbital lower than than of a3p orbital? Why is the energy of a 3p orbital lower

than the energy of a 3d orbital?

A qualitative explanation is found in the concept ofeffective nuclear charge “seen” by an electron

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Effective charge, Zeff, seen by valence electrons*

*Note x-axis is incorrect. What should it be?

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Two new features for the transition elements: the energy“switch” of the 4s and 3d orbitals and the Cr and Cu anomalies.

The 4s and 3d orbitals are very close in energy in the oneelectron atom.

Depending on the screening (Eeff) 4s < 3d or 3d < 4s.

From detailed computation and from experiment:4s < 3d in the neutral atoms at the beginning of the 4th period

20Ca 18[Ar]4s2 not [Ne]3s23p63d2

21Sc = 18[Ar]4s23d1:

Note in the cations of transition metals, 3d is lower than 4s:21Sc+2 = 18[Ar]3d1 not 18[Ar]4s1. Thus, 4s electrons are loss first!

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Cu: [Ar]4s23d9 is expected to be the ground state from Aufbaubut [Ar]4s13d10(a half filled subshell and a filled subshell) is theground state

Some rules about filling shell and subshells for n = 3 or greater:

There is a special stability of a filled subshell and a half-filled subshell.

This special stability of filled and half-filled subshells causes the Cr andCu electron ground state configurations to be different from thatpredicted from the Aufbau principle alone

Cr: [Ar]4s23d4 is expected to be the ground state from Aufbaubut [Ar]4s13d5 (two half filled subshells) is the ground state

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“Expected” Found21Sc 18[Ar]4s23d22Ti 18[Ar]4s23d2

23V 18[Ar]4s23d3

24Cr 18[Ar]4s23d4 18[Ar]4s1(↑)3d5 (↑↑↑↑↑)25Mn 18[Ar]4s23d5 half filled half filled26Fe 18[Ar]4s23d6

27Co 18[Ar]4s23d7

28Ni 18[Ar]4s23d8

29Cu 18[Ar]4s23d9 18[Ar]4s1 (↑)3d10 (↑↓↑↓↑↓↑↓↑↓)30Zn 18[Ar]4s23d10 half filled filled

The “surprises” for electron configurations at 24Cr and 29Cu are due tothe special stability of half filled subshells and filled subshells.

Following the (n + l) rule, the electron configurations of the transitionelements of the fourth row

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Ground State Electron Configurations: d block elements Paramagnetism and Hund’s Rule

Diamagnetic

Diamagnetic

Paramagnetic

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Row Configuration Shorthand

First: 1s2 2[He]

Second: 2s22p6 10[Ne]

Third: 3s23p6 18[Ar]

Fourth: 4s23d104p6 36[Kr]

Fifth: 5s24d105p6 54[Xe]

The closed shell electron configurations ofmultielectron atoms

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The electron configurations of the valence shell of themain group elements

Group I: ns1np0 The alkali metals: (H) Li, Na, K, Rb, Cs

Group II: ns2np0 The alkali earth metals: Be, Mg, Ca, Sr, Ba

Group III: ns2np1 The boron family: B, Al, Ga, In, Tl

Group IV: ns2np2 The carbon family: C, Si, Ge, Sn, PbGroup V: ns2np3 The nitrogen family: N, P, As, Sb, Bi

Group VI: ns2np4 The chalcogens: O, S, Se, Te, Po

Group VII: ns2np5 The halogens: F, Cl, Br, I

Group VIII: ns2np6 The noble gases: (He) Ne, Ar, Kr, Xe,

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The periodic table by orbital filling

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The atomic electron configurations of first five rows of theperiodic table give the elements their signature characteristics of

metals and non-metals

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Electronic structure and the periodic table

Electrons in the outermost shell of an atom are the mostimportant in determining chemical properties. Chemical reactions

involve only the outer (valence) electrons. The inner (core)electrons are not involved in chemical reactions.

Elements in a given vertical column (families) of the periodictable have similar outer-shell electron configurations andsimilar properties. They are isoelectronic with respect to

the number of valence electrons.

Elements in a row show regular trends in their properties dueto the continuing increase in the number of valence electrons

until a shell is filled.

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Summary: The Periodic Table built up by electron configurations: theground state electron configurations of the valence electrons of the

elements

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Effective nuclear charge (Zeff) on the outer electrons

Maintain hydrogen atom like orbitals as an approximation,but subshell energies are not equal: Ens < Enp < End < Enf

A s electron penetrates to the nucleus more than a pelectron: a p electron penetrates to the nucleus more thana d electron: more penetration, more stable, lower energy.

Subshell energies: E3s < E3p < E3d

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Classification of orbitals of a many electron atomaccording to their energies.

A group of orbitals with exactly equal energiescomprise a subshell.

Example: 2px, 2py and 2pz

Orbitals with same value of n and different valueof l comprise a shell.

Example: 2s and 2p comprise a shell.

The orbital approximation ignores electron-electronrepulsion, but takes into account Hund’s rule: electronswith parallel spins (↑↑) tend to stay apart compared to

electrons with antiparallel spins (↑↓).

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More about orbitals, shells and subshells

Each shell of principal quantum number n contains n subshellsn = 1, only one subshell (s)n = 2, two subshells (s, p)n = 3, three subshells (s, p, d)

Each subshell of quantum number l contains (2l + 1) orbitalsl = 0, (2x0 + 1) = 1 orbital (s)l = 1, (2x1 + 1) = 3 orbitals (px, py, pZ) l = 2, (2x2 + 1) = 5 orbitals (dxy, dyz, dxz, dx2 - y2, dz2)

The number of orbitals for a given n is n2 (solutions to wave equation)For n = 1, one orbital; for n = 2, four orbitals, for n = 3, nine orbitals

The number of electrons that can fill a given shell = 2n2

For each orbital, there can be a maximum of 2 occupying electrons

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The Pauli principle imposes structure on the manyelectron atom.

Without it, all the electrons might be expected tocrowd into the low energy orbitals. With it the

electrons are organized, filled orbitals with no morethan two electrons.

The ground state of an atom possesses the lowestenergy organization of electrons around the nucleus.All electron organizations are described by electron

configurations.

The ground state of an atom corresponds to thelowest energy electron configuration.

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Orbital shells and the buildingup of the periodic table

A shell is a set of orbitalswith the same value of n

The 18Ar atom(1s22s2263s23p6) has shells

as shown (left) in the profileof electron density as a

function of distance fromthe nucleus

The last shell (3s23p6) containsthe 8 valence electrons of our

Lewis structures!

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The Pauli exclusion principle (A principle not derived fromquantum mechanics, but required to explain observations).

Two equivalent statements of the Pauli exclusion principle:

(1) No two electrons may have the same set of four quantumnumbers;

(2) No more than two electrons may occupy the same orbital.

Because of the Pauli exclusion principle, outer electrons do not“fall” into the inner shell. Thus, the atom is stable.

Furthermore, the principle together with the Aufbau principleleads to “magic” number of electrons in shells

Building up electronic configurations of atoms:In Place: the ordering of the energy levels (n +l) rule.

Two principles now needed to be applied: The Pauli exclusionprinciple and the Aufbau principle


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