I. The Periodic Law and the Periodic Table configuration of all elements are constructed by filling lowest energy orbitals sequentially in what is called the “Aufbau Process”

Announcements--Exam 3 Oct 3..............Includes chapters 7/8/9/10The excluded items include:1. Classical distinction between energy and matter (p. 217)2. Numerical problems involving the Rydberg equation (equations

7.3 and 7.4)3. Spectral analysis in the laboratory (pp. 226-227)4. Numerical problems involving the Heisenberg uncertainty

principle (p. 231)5. Trends among the transition elements (p. 261)6. Trends in electron affinity (pp. 265-266)7. Pseudo-noble gas configuration (p. 269)8. Lattice energy (pp. 283-285)9. IR spectroscopy (p. 292)10.Numerical problems involving electronegativity (p. 296)11.Electronegativity and oxidation number (p. 297)12.Section 11.3: MO theory and electron delocalization13.All sections in chapter 12 except 12.3

is

having

amplitude energy frequency wavelength

related by

units calleda wave and a particle

related by

Quanta

EM radiation

are

emittedabsorbed

electrons

involve energy changes in

atomsatoms

moleculesdescribed by

Wave functions e- configuration comprising

determined by

Aufbau Rules

E = hv c = !v

e- fillinggives

having quantum uunmbers

Wave Function (Orbital)

described by

described by

which are

spdf electronic configuration

Aufbau Rules

determined by

which involve

comprising

Core Electrons

Valence Electrons

Periodic Table

basis for

which summarizes

Periodic Properties

Hund’s Rule

OribitalEnergy

PauliExclusion

e- filling

Quantum Numbers

Quantum Numbers

Principaln = 1,2,3,..

Angular momentum, l

Magneticml

Spin, ms

defines

defines

defines

defines

defines

Orbital size & energy

Electron spin

Orbitalorientation

Orbitalshape

Chapter 8

Electronic Configuration and Periodicity

I. The Periodic Law and the Periodic Table

• 1864 Newland “Law of Octaves

• 1869 Dimitri Mendeleev and Lother Meyer

When the elements are arranged in order of increasing atomic mass, certain sets of chemical and physical properties recur periodically.

• 1913 Henry Mosely relates X-ray frequency to atomic number

When the Elements Were Discovered

1. Principal Quantum Number (n): Defines the size and energy level of the orbital. n = {1,2,3,4,.....}. Also called a shell (K = 1, L = 2, M = 3, N = 4, .....).

2. Angular Momemtum Quantum Number (l): Defines the “shape” of the orbital which is a volume in space where the electron is likely to be found. Also called a subshell. l = {0,1,2,3...up to n-1} where (0=s, 1=p, 2=d, 3=f)

3. Magnetic Quantum Number (ml): Defines the spatial orientation of an orbital of the same energy. ml = {-l, 0, +l}

4. Magnetic Spin Quantum Number (ms): Defines the orientation of “electron spin”. ms = {+1/2 or -1/2}.

Quantum numbers (n,l,ml,ms) specify “allowed states” or “orbitals” which are regions of space where electrons are likely to be found around the nucleus.

Electronic configuration of the elements: four quantum numbers describe an electron in a ground state atom.

Name

principal

Symbol Permitted Values Property

n positive integers (1,2,3,!)

orbital energy (size)

angular momentum

l integers from 0 to n-1 orbital shape

magnetic ml integers from -l to 0 to +l

orbital orientation in space

spin ms +1/2 or -1/2 direction of e- spin

The lowest energy (ground state) electronic configuration of all elements are constructed by filling lowest energy orbitals sequentially in what is called the “Aufbau Process”.

1. Lower energy (n-quantum number) orbitals fill first.

2. Hund’s Rule-orbitals fill one electron at a time before electrons are paired.

3. Pauli Exclusion Principle: No two electrons can have same 4-quantum numbers)

Electrons fill the lowest energy orbitals first, 2 at a time! 1s

2s

2p

3s3p4s

The order of filling of the orbitals can be remembered using a mnemonic device. Memorize this to help you!

For an Hydrogen atomorbital energy only dependson the n quantum number.

For many electrons atomsthe energy of an orbital or electron depends on bothn and l (3s < 3p < 3d)

Chemists use spdf notation and orbital box diagrams to denote or show the “ground state electronic configuration” of elements.

spdf Notation

orbital box diagram

H

He

1s1

1s2

Element

n principal quantum #

l quantum number # of electronsin orbital

Spin quantum number. An arrow denotes an electron with “spin up” (+1/2) or “spin-down” (-1/2).

The Pauli Exclusion principle states: “No two electrons can have the same 4-quantum numbers”. The spin numbers can not be the same (spin up and spin down allowed only).

Li 1s22s1

Example:

(n, l, ml and ms)

Atomic Number/Element

Orbital BoxDiagram

Full-electronicconfiguration

Condensed-electronicconfiguration

[He]2s1

The order of filling of the orbitals can be remembered using a mnemonic device. Memorize how to write it out as it determines electronic structure.

s-blockp-block

d-block

f-block

transition metals

inner transition metals

main groupmain group

H

He

Li

Be

1s1

1s2

1s22s2

1s22s1


Orbital BoxDiagram



1s1

1s2

[He]2s1

[He]2s2

written with noble gas configuration

Electronic configuration using Aufbau Process

1s22s22p3

1s22s22p4

1s22s22p5

1s22s22p6

1s22s22p1

1s22s22p2

B

C


Orbital BoxDiagram



[He]2s22p1

[He]2s22p2

[He]2s22p3

[He]2s22p4

[He]2s22p5

[He]2s22p6

Odd-filling behavior here!4th and 9th position.

Paramagnetic

unpaired electrons

2p

Diamagnetic

all electrons paired

2p

• Diamagnetic atoms or ions:– All e- are paired.– Weakly repelled in a magnetic field.

• Paramagnetic atoms or ions:– Unpaired e- exist in an orbital– Attracted to an external magnetic field.

Unpaired electrons in orbitals gives rise to paramagnetism and is attracted to a magnetic field. Diamagnetic species contain all paired electrons and is “repelled” by the magnetic field.

Paramagnetic Diamagnetic

Unpaired electrons in orbitals gives rise to paramagnetism and is attracted to a magnetic field. Diamagnetic species contain all paired electrons and is “repelled” by the magnetic field.

Magnetic field off Magnetic field on Magnetic field on

Paramagentic Diamagentic

Fe: [Ar]4s23d6 Fe2+: [Ar]4s03d6 or [Ar]3d6

Mn: [Ar]4s23d5 Mn2+: [Ar]4s03d5 or [Ar]3d5

Fe3+: [Ar]4s03d5 or [Ar]3d5Fe: [Ar]4s23d6

When a cation is formed from an atom of a transition metal, electrons are removed first from the ns orbital, then from the (n-1)d orbital.

Metals lose electrons so that cation has a noble-gas outer electron configuration.

Na [Ne]3s1 Na+ [Ne]

Ca [Ar]4s2 Ca2+ [Ar]

Al [Ne]3s23p1 Al3+ [Ne]

Non-metals gain electrons so that anion has a noble-gas outer electron configuration.

H 1s1 H- 1s2 or [He]

F 1s22s22p5 F- 1s22s22p6 or [Ne]

O 1s22s22p4 O2- 1s22s22p6 or [Ne]

N 1s22s22p3 N3- 1s22s22p6 or [Ne]

Metals loose electrons (oxidized) to become cations. Non-metals gain electrons to become anions. The electronic configuration of each reflects this change in the number of electrons.

Na+, Al3+, F-, O2-, and N3- are all said to be “isoelectronic with Ne” as they have the same electronic configuration....all subshells are filled.

Isoelectronic species are two different elements with the same electronic configuration--but not the same nuclear configuration.

Na: [1s22s22p63s1] =====> Na+: [1s22s22p6] = [Ne]oxidation

oxidationAl: [1s22s22p63s23p1] =====> Al3+: [1s22s22p6] = [Ne]

N: [1s22s22p3] =====> N3-: [1s22s22p6] = [Ne]reduced

O: [1s22s22p4] =====> O2-: [1s22s22p6] = [Ne]reduced

F: [1s22s22p5] =====> F-: [1s22s22p6] = [Ne]reduced

Metals and non-metals form ions with electronic configurations closest to their nearest noble gas configuration.

3A 4A 5A 6A 7A 8A2A1A

Metals and non-metal ions tend to form electronic states closest to their nearest noble gas configuration.

What is the spdf and condensed electron configuration of Mg and Mg2+ ? Mg 12 electrons

What are the possible quantum numbers for the last (outermost) electron in Cl?

Cl 17 electrons 1s < 2s < 2p < 3s < 3p < 4s 1s22s22p63s23p5 2 + 2 + 6 + 2 + 5 = 17 electronsLast electron added to 3p orbitaln = 3 l = 1 ml = -1, 0, or +1 ms = ! or -!

C) Is ground state F paramagenetic or diamagnetic?

Mg 1s22s22p63s2 [Ne]3s2

Mg2+ 1s22s22p63s0 [Ne]3s0 = [Ne]

Using the periodic table on the inside cover of the text and give the full and condensed electrons configurations, partial orbital diagrams showing valence electrons, and number of inner electrons for the following elements:

(a) potassium (K: Z = 19)

(b) molybdenum (Mo: Z = 42)

(c) lead (Pb: Z = 82)

(b) for Mo (Z = 42) 36 inner electrons and 6 valence electrons1s22s22p63s23p64s23d104p65s14d5

[Kr] 5s14d5

(c) for Pb (Z = 82) 78 inner electrons and 4 valence electrons.

[Xe] 6s24f145d106p2

condensedpartial orbital diagram

full configuration

condensed

partial orbital diagram

full configuration 1s22s22p63s23p64s23d104p65s24d10

5p66s24f145d106p2

6s2 6p2

5s1 4d5 5p

(a) for K (Z = 19)1s22s22p63s23p64s1

[Ar] 4s1

4s1

condensedorbital diagram

full configurationThere are 18 inner electrons.

3d 4p

Use condensed electron configurations to write the reaction for the formation of each transition metal ion, and predict whether the ion is paramagnetic.

(a) Mn2+(Z = 25) (b) Cr3+(Z = 24) (c) Hg2+(Z = 80)

Write the electron configuration and remove electrons starting with ns to match the charge on the ion. If the remaining configuration has unpaired electrons, it is paramagnetic.

SOLUTION:

paramagnetic(a) Mn2+(Z = 25) Mn([Ar]4s23d5) Mn2+ ([Ar] 3d5) + 2e-

(b) Cr3+(Z = 24) Cr([Ar]4s13d5) Cr3+ ([Ar] 3d3) + 3e-

paramagnetic

(c) Hg2+(Z = 80) Hg([Xe]6s24f145d10) Hg2+ ([Xe] 4f145d10) + 2e-

not paramagnetic (is diamagnetic)

Use condensed electron configurations to write the reaction for the formation of each transition metal ion, and predict whether the ion is paramagnetic.

(a) Mn2+(Z = 25) (b) Cr3+(Z = 24)

Write the electron configuration and remove electrons starting with ns to match the charge on the ion. If the remaining configuration has unpaired electrons, it is paramagnetic.

(c) Hg2+(Z = 80)

Identify n and l quantum numbers for each of the following.

What neutral element has the following orbital-filling diagram?

fourth shellthird shell

Identify n and l quantum numbers for each of the following.

What neutral element has the following orbital-filling diagram?

fourth shellthird shell

3p 4dz2

Gallium = Ga

Many atomic properties show periodicity and trends.

Amount of energy to remove 1 mole e- from 1 mole of gaseous atoms or element

Amount of energy to add 1 mole e- to 1 mole of gaseous atoms or element


1) Inner core electrons : electrons filling the lower n shells of an element. They are located closer to the nucleus.

2. Outer core or VALENCE e- : those e- in the highest energy level (highest n-value). The number of valence e- is given by the Group Number in the periodic table for Group A.

--Responsible for chemistry and bonding of elements forming compounds or ions (true for representative but not transition metals--more complex).

Electrons in elements are categorized either as inner core electrons or valence electrons.

Periodicity in the chemical reactivity of elements occurs because of periodicity in the electronic structure of valence electrons!

1-electron outer s-orbital

2-electrons outer d-orbital

5-electrons outer p-orbital

Inner core electrons “shield” outer electrons from the positive charge of the nucleus.

Effective nuclear charge (Zeff) is the electrostatic force felt by the outer valence electrons taking into “shielding” by core electrons.

To a good approximation: effective nuclear charge, Zeff is given by:

Zeff = Z – core e-

# of inner non-valence electrons

# protonsEffective Nuclear charge

Bigger Zeff means more “pull” or electrostatic force between nucleus and electrons.

Configuration Element Z (p+)Core

ElectronsValence

Electrons ZeffectiveRadius(pm)

[Ne]3s1 Na 11 10 1 1 186[Ne]3s2 Mg 12 10 2 2 160

[Ne]3s23p1 Al 13 10 3 3 143

[Ne]3s23p2 Si 14 10 4 4 132

[Ne]3s23p3 P 15 10 5 5 128

[Ne]3s23p4 S 16 10 6 6 127

[Ne]3s23p5 Cl 17 10 7 7 99

[Ne]3s23p6 Ar 18 10 8 8 98

[Ar]4s1 K 19 18 1 1 227

[Ar]4s2 Ca 20 18 2 2 197

[Ar]4s23d1 Sc 21 18 3 3 135

Zeff = Z – core e-****

Because of increasing effective nuclear charge across a period, atomic radii decrease across a Period. As n increases down a group so does the radius.

Incr

easi

ng A

tom

ic R

adiu

sDecreasing Atomic Radius

n increases


Amount of energy to remove 1 mole e- from 1 mole of gaseous atoms or element

Amount of energy to add 1 mole e- to 1 mole of gaseous atoms or element

Periodicity of Atomic Radius

Group I

Group VIII

Cations get smaller (greater Zeff) Anions get larger (lower Zeff)

The radii of cations are smaller than their parent neutral atoms, while anions are larger than its parent.

Using only the periodic table rank each set of main group elements in order of decreasing atomic size:

(a) Ca, Mg, Sr

(b) K, Ga, Ca

(c) Br, Rb, Kr

(d) Sr, Ca, Rb

Using only the periodic table rank each set of main group elements in order of decreasing atomic size:

(a) Ca, Mg, Sr

(b) K, Ga, Ca

(c) Br, Rb, Kr

(d) Sr, Ca, Rb

SOLUTION:(a) Sr > Ca > Mg These elements are in Group 2A(2).

(b) K > Ca > Ga These elements are in Period 4.

(c) Rb > Br > KrRb has a higher n engery level and is far to the left. Br is to the left of Kr.

(d) Rb > Sr > Ca Ca is one energy level smaller than Rb and Sr. Rb is to the left of Sr.

First ionization energies of the main-group elements. First Ionization Energy

Ionization energy is the minimum energy (kJ/mol) required to remove an 1 mole of electrons from one mole of a gaseous atom in its ground state (!H > 0).

I1 + X (g) X+(g) + e-

I2 + X (g) X2+(g) + e-

I3 + X (g) X3+(g) + e-

I1 first ionization energy

I2 second ionization energy

I3 third ionization energy

I1 < I2 < I3

Ranking Elements by First Ionization Energy

PLAN:

PROBLEM: Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE1:

(a) Kr, He, Ar (b) Sb, Te, Sn (c) K, Ca, Rb (d) I, Xe, Cs

IE decreases as you proceed down in a group; IE increases as you go across a period.



PLAN:

SOLUTION:

PROBLEM: Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE1:

(a) Kr, He, Ar (b) Sb, Te, Sn (c) K, Ca, Rb (d) I, Xe, Cs

IE decreases as you proceed down in a group; IE increases as you go across a period.

(a) He > Ar > Kr

(b) Te > Sb > Sn

(c) Ca > K > Rb

(d) Xe > I > Cs

Group 8A(18) - IE decreases down a group.

Period 5 elements - IE increases across a period.

Ca is to the right of K; Rb is below K.

I is to the left of Xe; Cs is furtther to the left and down one period.

1s2 2s2 2p3

1s2 2s2 2p1

1s2 2s2

1s2 2s1

1s2 2s2 2p2

1s2 2s2 2p4

The ionization energy increases dramatically when an core electron is removed from a non-valence shell.

Identifying an Element from Successive Ionization Energies

PLAN:

Name the Period 3 element with the following ionization energies (in kJ/mol) and write its electron configuration:

IE1 IE2 IE3 IE4 IE5 IE6

1012 1903 2910 4956 6278 22,230

Look for a large increase in energy which indicates that all of the valence electrons have been removed.The number valence electrons is reflected in the periodic table for Group A elements....find the group with that number of valence electrons.

IE6

22,230

SOLUTION:

The largest increase occurs after IE5, that is, after the 5th valence electron has been removed. Five electrons would mean that the valence configuration is 3s23p3 and the element must be phosphorous, P (Z = 15).The complete electron configuration is 1s22s22p63s23p3.

Identifying an Element from Successive Ionization EnergiesName the Period 3 element with the following ionization energies (in kJ/mol) and write its electron configuration:

IE1 IE2 IE3 IE4 IE5

1012 1903 2910 4956 6278

Main Group (or representative) metals form ionic basic oxides when reacted with oxygen while non-metals form covalent acidic oxides with oxygen.

Increasing Acidity

Li2O BeO B2O3 CO2 OF2

K2O

Cs2O

Rb2O

CaO

SrO

BaO

Ga2O3

Tl2O3

In2O3

GeO2

PbO2

SnO2

SeO3

TeO3

Br2O7

I2O7

6A(16)

7A(17)

4A(14)

3A2A1A

2

4

5

6

Increasing Basicity Na2O

IonicOxides

CovalentOxides basic acidic

Properties of Oxides Across a Period

Documents

I. The Periodic Law and the Periodic Table configuration of all elements are constructed by filling lowest energy orbitals sequentially in what is called the “Aufbau Process”