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Chemistry 4362. Advanced Inorganic Chemistry. Instructor: Dr. Byron K. Christmas Class Time: Tue & Thur - 5:29 to 6:49 p.m. Classroom: N-936  C-320 Phone: (713) 221-8169 FAX: (713) 221-8528 E-Mail: ChristmasB@uhd.edu. Introduction.

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chemistry 4362

Chemistry 4362

Advanced Inorganic Chemistry

Instructor: Dr. Byron K. Christmas

Class Time: Tue & Thur - 5:29 to 6:49 p.m.

Classroom: N-936  C-320

Phone: (713) 221-8169 FAX: (713) 221-8528

E-Mail: ChristmasB@uhd.edu

introduction
Introduction
  • “The chemistry of everything that is NOT organic…”

What is Inorganic Chemistry ?

  • “The chemistry of all of the elements and their compounds
  • except for the hydrocarbons and their derivatives.”
  • “The branch of chemistry falling between and overlapping
  • with physical chemistry and organic chemistry.”
  • “What Inorganic Chemists Do!”
  • Your Personal Definition??
what do inorganic chemists do
What Do Inorganic Chemists Do ?
  • Synthesize and characterize substances other than those that are clearly “organic”.
  • Determine the structures of inorganic substances.
  • Investigate the chemical reactions of inorganic substances.
  • Investigate the physical properties of inorganic substances.

.

  • Develop hypotheses and theories to explain and systematize
  • the empirical data collected.
why should you study inorganic chemistry
Why Should You Study Inorganic Chemistry ?

Elemental Composition of the Sun and the Universe

Sun Universe

Hydrogen92.5 % 90.87 %

Helium7.3 % 9.08 %

All Others 0.2 % 0.05 %

  • Essentially the entire universe is Inorganic.
  • The Earth is predominantly Inorganic.

Elemental Composition of the Earth’s Crust

Oxygen 45.5 % Iron 6.20 %

Silicon 27.2 % Calcium 4.66 % Aluminum 8.30 % All Others 8.14 %

inorganic materials are an essential part of our national economy
Inorganic materials are an essential part of our national economy.

U.S. Production of Top 10 Chemicals (x 109 lb.) - 1997

Sulfuric Acid 95.58

Nitrogen 82.88*

Oxygen 64.84*

Ethylene 51.08

Lime 42.56

Ammonia 38.39

Propylene 27.53

Phosphoric Acid 26.83

Ethylene Dichloride 26.29

Sulfur 26.24

From C&EN, June 29, 1998

*Calculated from “billion cubic feet at STP”

slide7

U.S. Production of Top 50 Chemicals (x 109 lb.) - 1994

Total Organics 279.17

Total Inorganics 450.19

Grand Total 729.36

  • Inorganics are essential to life.
    • Water is essential for all life.
    • About 30 different elements are believed to be

essential to life - 28 in addition to carbon and

hydrogen.

  • For all practical purposes, Inorganic Chemistry IS
  • chemistry - the study of the properties, composition, and structure of matter, the physical and chemical changes it undergoes, and the energy liberated or absorbed during those changes.
slide8

Approaches to the Study of Inorganic Chemistry

  • Empirical Approach (Descriptive Chemistry)
    • Historically this was the way it was taught.
    • It involves essentially all memorization.
    • It is necessary for a complete understanding of Chemistry.
  • Theoretical Approach
    • It provides a framework for understanding the “why” of

descriptive chemistry.

    • It can provide “intellectual satisfaction”.
    • It is limited in its ability to give explanations for all

observed phenomena.

    • It has dominated the teaching of Inorganic Chemistry

for over 30 years.

introduction to descriptive chemistry
Introduction to Descriptive Chemistry

Definition: “…the study of the composition, structure,

and properties of matter….” - thefacts about the

elements and their compounds.

Comments from the Experts: “…evident differences

between this and previous editions…is the absence of much

theoretical material previously included…the continuing

rapid growth of chemistry…required the addition of impor-

tant new factsto all of the descriptive material…over the

years, become less persuaded of the value of certain types of

theorizing….Thus, we felt obliged to make space forfacts at

the expense of theoretical material.” Cotton and Wilkinson,

Advanced Inorganic Chemistry, 5th Edition, 1988.

slide10

Comments from the Experts: “The facts concerning

the properties and reactions of substances are the very

essence of chemistry. Facts undergo little if any

change in contrast to constantly changing theories.

Moreover, …a chemist needs a solid background of

facts in order to appreciate the need for theories….”

R. J. Gillespie in the Forward of Chemistry of the

Elements, Greenwood and Earnshaw, 1st Edition,

1984.

slide11

“Over the years, the theoretical part tended to grow at the expense of

the descriptive material….The theoretical part tended to become the

end rather than the means….By the 1970’s many teachers had to

abandon any attempt to cover descriptive inorganic chemistry in the

traditional sense. Thus we can encounter the student who can

write an erudite account of structural minutiae in copper(II)

chemistry, ligand field spectra and…,but who knows little about

the more mundane compounds of the transition elements and

would be hard pressed to locate indium in the Periodic Table, let

alone venture anything about its chemistry.” Derek W. Smith,

Inorganic Substances: A Prelude to the Study of Descriptive

Inorganic Chemistry, 1st Edition, 1990.

slide12

“Chemistry has always been, and still is, a practical subject….An

American professor told me he divided inorganic chemistry books into

two types: theoretical and practical. In deciding how to classify any

particular book, he first looked to see if the extraction of the two most

produced metals (Fe and Al) was adequately covered, what impurities

were likely to be present, and how the processing was adapted to re-

move them. Second, he looked to see if the treatment of the bonding

in xenon compounds and ferrocene was longer than that of the pro-

duction of ammonia. Third, he looked to see if the production and

uses of phosphates were covered adequately….For some years there

has been a trend for chemistry teaching to become more theoretical.

There is always theoretical interest in another interesting oxidation

state or another unusual complex, but the balance of this book is

tilted to ensure that they do not exclude the commonplace, the mun-

dane and the commercially important.”J. D. Lee, Concise Inorganic

Chemistry, 5th Edition,1996.

slide13

Industrial Applications Approach

    • Few schools other than chemical engineering

programs have used this approach.

    • It is of great “practical” importance for students

preparing for industry.

    • It is of limited utility in preparing for graduate

work in chemistry.

  • Balanced Approach
    • Provides a balance among all approaches.
    • Applicable to “survey-type” course.
    • Useful for either graduate school or industry

preparation.

    • Used in THIS COURSE!!
course overview
Course Overview
  • Introduction to Inorganic Polymers
  • Theoretical Concepts
    • Atomic Structure & the Periodic Table
    • Properties of the Elements
    • Introduction to Chemical Bonding
    • The Covalent Bond
    • The Metallic Bond
    • The Ionic Bond
    • Intermolecular Attractive Forces
    • Inorganic Thermodynamics and Kinetics
    • Solvent Systems and Acids and Bases
    • Oxidation/Reduction
  • Descriptive Chemistry and Industrial Applications
  • Student Presentations
slide15

Introduction to Inorganic Polymers

(Page 338 in Text)

Questions to Ponder:

1. Would you know an Inorganic Polymer if you saw one?

2. How could you determine if an inorganic material was, in fact,

“polymeric”?

3. List important types of Inorganic Polymers.

4. How would you determine what is and what is NOT an

Inorganic Polymer?

Is NaCl a Polymer? Is Graphite a Polymer? What about Diamond?

Is Aluminum a Polymer? What about Window Glass?

5. What general principles of chemical bonding, atomic size, etc.

lead to effective polymer formation for different types of

elements?

6. What are commercially important inorganic polymers?

slide16

Introduction to Inorganic Polymers

Catenation – What are the requirements?

Valence of two or more?

Bond energies?

What else?

Homocatenation

Heterocatenation

Bond Energies – kJ/mole

C-C 356

Si-Si 222

Ge-Ge 188

Sn-Sn 167

Pb-Pb 87

S-S 251

P-P 201

O-O 142

Si-O 460

Sn-O 544

Al-O 586

Si-N 355

B-N 460

http://chemviz.ncsa.uiuc.edu/content/doc-resources-bond.html

slide17

Introduction to Inorganic Polymers

Polysulfide Demonstration/Experiment

Questions?

Assignments!

Study Hand-outs and your text on Inorganic Polymers

Find three to five ADDITIONAL references on the Web

and study them

Prepare for next Thursday’s Silicone laboratory

(Page 176 in Lab Manual)

theoretical concepts

Theoretical Concepts

Chapter 1

Atomic Structure & the Periodic Table

Properties of the Elements

Introduction to Chemical Bonding

The Ionic Bond

The Covalent Bond

The Metallic Bond

Intermolecular Attractive Forces

Thermodynamics

Acids and Bases

Oxidation/Reduction

slide19

ATOMIC STRUCTURE

Definition of Chemistry:

The study of the properties, composition, and

STRUCTURE of matter, the physical and

chemical changes it undergoes, and the energy

liberated or absorbed during those changes.

The foundation for theSTRUCTUREof inorganic

materials is found in theSTRUCTUREof the atom.

Material Properties

Bulk Structure

Molecular Structure

Atomic Structure

atomic structure
ATOMIC STRUCTURE

Historical Development:

  • Greek Concepts of Matter
    • Aristotle - Matter is continuous, infinitely
    • divisible, and is composed of only 4 elements:
    • Earth, Air, Fire, and Water
    • Won the philosophical/political battle.
    • Dominated Western Thought for Centuries.
    • Seemed very “logical”.
    • Was totally WRONG!!
atomic structure21
ATOMIC STRUCTURE

The “Atomists” (Democritus, Lucippus,

Epicurus, et. al.) - Matter consists ultimately

of “indivisible” particles called “atomos” that

canNOT be further subdivided or simplified.

If these “atoms” had space between them,

nothing was in that space - the “void”.

  • Lost the philosophical/political battle.
  • Lost to Western Thought until 1417.
  • Incapable of being tested or verified.
  • Believed the “four elements” consisted of

“transmutable” atoms.

  • Was a far more accurate, though quite imperfect

“picture” of reality.

atomic structure22
ATOMIC STRUCTURE

Modern Concepts of Matter

John Dalton (1803)- An atomist who formalized

the idea of the atom into a viable scientific theory

in order to explain a large amount of empirical

data that could not be explained otherwise.

  • Matter is composed of small “indivisible”particles

called “atoms”.

  • The atoms of each element are identical to each

other in mass but different from the atoms of other

elements.

  • A compound contains atoms of two or more

elements bound together in fixed proportions

by mass.

atomic structure23
ATOMIC STRUCTURE
  • A chemical reaction involves a rearrangement of

of atoms but atoms are not created nor destroyed

during such reactions.

Present Concepts - An atom is an electrically

neutral entity consisting of negatively charged

electrons (e-) situated outside of a dense, posi-

tively charged nucleus consisting of positively

charged protons (p+) and neutral neutrons (n0).

ParticleChargeMass

Electron - 1 9.109 x 10 -28 g

Proton +1 1.673 x 10 -24 g

Neutron 0 1.675 x 10 -24 g

atomic structure24
ATOMIC STRUCTURE

Nucleus

Model of a

Helium-4

(4He) atom

p+no

e-

e-

no p+

Electron Cloud

How did we get this concept? - This portion of our

program is brought to you by:

Democritus, Dalton, Thompson, Planck, Einstein, Millikan,

Rutherford, Bohr, de Broglie, Heisenberg, Schrödinger,

Chadwick, and many others.

atomic structure25
ATOMIC STRUCTURE
  • Democritus - First atomic ideas
  • Dalton - 1803 - First Atomic Theory
  • J. J. Thompson - 1890s - Measured the charge/mass
    • ratio of the electron (Cathode Rays)

Fluorescent

Material

_

Cathode

+

Anode

Electric Field

Source (Off)

With the electric field off, the cathode ray is not deflected.

atomic structure26
ATOMIC STRUCTURE

-

Fluorescent

Material

-

Cathode

+

+

Anode

Electric Field

Source (On)

With the electric field on, the cathode ray is deflected

away from the negative plate. The stronger the electric

field, the greater the amount of deflection.

-

Cathode

+

Anode

Magnet

atomic structure27
ATOMIC STRUCTURE

With the magnetic field present, the cathode ray is

deflected out of the magnetic field. The stronger the

magnetic field, the greater the amount of deflection.

e/m = E/H2r

e = the charge on the electron

m = the mass of the electron

E = the electric field strength

H = the magnetic field strength

r = the radius of curvature of the electron beam

Thompson, thus, measured the charge/mass ratio

of the electron - 1.759 x 108 C/g

atomic structure28
ATOMIC STRUCTURE
  • Summary of Thompson’s Findings:
  • Cathode rays had the same properties no matter

what metal was being used.

  • Cathode rays appeared to be a constituent of all

matter and, thus, appeared to be a “sub-atomic”

particle.

  • Cathode rays had a negative charge.
  • Cathode rays have a charge-to-mass ratio

of 1.7588 x 108 C/g.

atomic structure29
ATOMIC STRUCTURE

R. A. Millikan - Measured the charge of the electron.

In his famous “oil-drop” experiment, Millikan was able to

determine the charge on the electron independently of its

mass. Then using Thompson’s charge-to-mass ratio, he

was able to calculate the mass of the electron.

e = 1.602 10 x 10-19 coulomb

e/m = 1.7588 x 108 coulomb/gram

m = 9.1091 x 10-28 gram

Goldstein - Conducted “positive” ray experiments that

lead to the identification of the proton. The charge

was found to be identical to that of the electron and

the mass was found to be 1.6726 x 10-24 g.

atomic structure30
ATOMIC STRUCTURE

Ernest Rutherford - Developed the “nuclear” model

of the atom.

The Plum Pudding Model of the atom:

+

+

+

A smeared out “pudding”

of positive charge with

negative electron “plums”

imbedded in it.

- - - - -

- - - - -

- - - - -

+

+

+

+

+

+

+

The Metal Foil Experiments:

Fluorescent

Screen

a-particles

Radioactive

Material in

Pb box.

Metal

Foil

atomic structure31
ATOMIC STRUCTURE

If the plum pudding model is correct, then all of

the massive a-particles should pass right through

without being deflected.

In fact, most of the a - particles DID pass right

through. However, a few of them were deflected at

high angles, disproving the “plum pudding” model.

Rutherford concluded from this that the atom con-

sisted of a very dense nucleus containing all of the

positive charge and most of the mass surrounded by

electrons that orbited around the nucleus much as

the planets orbit around the sun.

atomic structure32
ATOMIC STRUCTURE

Assignment:

Assume the diameter of the nucleus of a hydrogen

atom is 1 x 10 -13 cm and the diameter of the atom

is 1 x 10 -8 cm.

1. Calculate the volume of the nucleus and the volume

of the atom in cm3 .

2. Calculate the volume of empty space in the atom.

3. Calculate the ratio of the volume of the nucleus to

volume of the whole atom.

4. Calculate the density of the nucleus if the proton’s

mass is1.6726 x 10-24 g

atomic structure33

Problems with the Rutherford Model:

It was known from experiment and electromagnetic

theory that when charges are accelerated, they

continuously emit radiation, i.e., they loose energy

continuously. The “orbiting” electrons in the atom

were, obviously, not doing this.

ATOMIC STRUCTURE

Planck

  • Atomic spectra and blackbody radiation

were known to be DIScontinuous.

Bohr

  • The atoms were NOT collapsing.
atomic structure34
ATOMIC STRUCTURE

Atomic Spectra - Since the 19th century, it had

been known that when elements and compounds

are heated until they emit light (glow) they emit

that light only at discrete frequencies, giving a

line spectrum.

-

+

Hydrogen

Gas

Line Spectrum

atomic structure35
ATOMIC STRUCTURE

When white light is passed through a sample of

the vapor of a substance, only discrete frequencies

are absorbed, giving an absorption ban spectrum.

These frequencies are identical to those of the

line spectrum of the same element or compound.

For hydrogen, the spectroscopists of the 19th

Century found that the lines were related by the

Rydberg equation:

n/c = R[(1/m2) - (1/n2)]

n = frequency

R = Rydberg Constant

c = speed of light

m = 1, 2, 3, ….

n = (m+1), (m+2), (m+3), ….

atomic structure36
ATOMIC STRUCTURE

Max Planck - In 1900 he was investigating the nature

of black body radiation and tried to interpret his

findings using accepted theories of electromagnetic

radiation (light). He was NOT successful since these

theories were based on the assumption that light had

WAVE characteristics.

To solve the problem he postulated that light was

emitted from black bodies in discrete packets he

called “quanta”. Einstein later called them

“photons”. By assuming that the atoms of the black

body emitted energy only at discrete frequencies, he

was able to explain black body radiation.

E = hn

atomic structure37

Both spectroscopy and black body radiation

indicated that atoms emitted energy only at

discrete frequencies or energies rather than

continuously.

ATOMIC STRUCTURE

Is light a particle or a wave??

Why do atoms emit only discrete energies?

What actually happens when light interacts

with matter?

What was wrong with Rutherford’s Model?

atomic structure38

Niels Bohr - Bohr corrected Rutherford’s model

  • of the atom by formulating the following postulates:
  • Electrons in atoms move only in discrete orbits

around the nucleus.

  • When in an orbit, the electron does NOT emit

energy.

  • They may move from one orbit to another but are

NEVER residing in between orbits.

  • When an electron moves from one orbit to

another, it absorbs or emits a photon of light with a

specific energy that depends on the difference in

energy between the two orbits.

ATOMIC STRUCTURE
atomic structure39
ATOMIC STRUCTURE

Balmer

Series

(Visible)

Lyman

Series

+

Paschen

Series

(UV)

(IR)

The Bohr Model of the Atom

atomic structure40
ATOMIC STRUCTURE
  • The lowest possible energy state for an electron

is called the GROUND STATE. All other states

are called EXCITED STATES.

En = (- 2.179 x 10-18 J)/n2

Ephoton = Ehigh - Elow

Ephoton = [(- 2.179 x 10-18 J)/n2high]

-[(- 2.179 x 10-18 J)/n2low]

= - 2.179 x 10-18 J[(1/n2high) - (1/n2low)]

Does this equation look familiar?

n/c = R[(1/m2) - (1/n2)]

atomic structure41
ATOMIC STRUCTURE

Niels Bohr won the Nobel Prize for his work.

However, the model only worked perfectly for

hydrogen. What about all of those other elements??

Louis de Broglie - Thought that if light, which was

thought to have wave characteristics, could also have

particle characteristics, then perhaps electrons, which

were thought to be particles, could have characteristics

of waves.

l = h/mv

An electron in an atom was a “standing wave”!

atomic structure42
ATOMIC STRUCTURE

Werner Heisenberg - Developed the “uncertainty”

principle: It is impossible to make simultaneous and

exact measurements of both the position (location)

and the momentum of a sub-atomic particle such as

an electron.

(Dx)(Dp)  h/2p

Our knowledge of the inner workings of atoms and

molecules must be based on probabilities rather

than on absolute certainties.

atomic structure43
ATOMIC STRUCTURE

Erwin Schrödinger - Developed a form of quantum

mechanics known as “wave mechanics”.

Hy = Ey

H = Hamiltonian operator

E = Total energy of the system

y = Wave function

[(-h2)/(8p2m)]2 - [e2/r]  = E

Kinetic Energy

Term

Potential Energy

Term

This is simply a quantum mechanical statement of the Law

of Conservation of Energy

atomic structure44

Of the numerous solutions to the Schrödinger equation

  • for hydrogen, only certain ones are allowed due to the
  • following boundary conditions:
  • Y, the wave function, must be continuous and finite.
  • It must be single-valued at all points (There can’t be

two different probabilities of finding an electron at one

point in space).

  • The probability of finding the electron, Y2, somewhere

in space must = 1.

ATOMIC STRUCTURE

+



Y2dxdydz = 1

- 

Y has many values that meet these conditions. They are

called “orbitals”.

atomic structure45
ATOMIC STRUCTURE
  • Wave Function - A mathematical function associated
  • with each possible state of an electron in an atom or
  • molecule.
    • It can be used to calculate the energy of an
    • electron in the state
    • the average and most probable distance from the

nucleus

    • the probability of finding the electron in any

specified region of space.

Y

Y

Y

Y

Y

atomic structure46
ATOMIC STRUCTURE

Quantum Numbers:

Principle Quantum Number, n - An integer

greater than zero that represents the principle

energy level or “shell” that an electron occupies.

Energy # of orbitals

n Level Shell n2

1 1st K 1

2 2nd L 2

3 3rd M 9

4 4th N 16

etc. etc. etc. etc.

atomic structure47
ATOMIC STRUCTURE

Azimuthal Quantum Number, l - The quantum

number that designates the “subshell” an electron

occupies. It is an indicator of the shape of an orbital

in the subshell. It has integer values from 0 to n-1.

l = 0, 1, 2, 3, …, n - 1

s p d f

Magnetic Quantum Number, ml - The quantum

number that determines the behavior of an electron

in a magnetic field. It designates the orbitaland

has integer values from -l to +l including 0.

ml = -l, …, -3, -2, -1, 0, +1, +2, +3, …, +l

atomic structure48

Orbital # of

n l Name ml Orbitals

1 0 1s 0 1

2 0 2s 0 1

1 2p -1, 0, +1 3

3 0 3s 0 1

1 3p -1, 0, +1 3

2 3d -2, -1, 0, +1, +2 5

etc. etc. etc. etc. etc.

ATOMIC STRUCTURE

Spin Quantum Number, ms - The quantum number

that designates the orientation of an electron in a

magnetic field. It has half-integer values, +½ or -½.

atomic structure49
ATOMIC STRUCTURE

So what do atoms look like?

A. Interpretation of Y: Theprobability of finding

an electron in a small volume of space centered

around some point is proportional to the value of

Y2at that point.

B. Electron Probability Density vs. r

C. Dot Density Representation: Imagine super-

imposing millions of photographs taken of an

electron in rapid succession.

D. Radial Densities

atomic structure50
ATOMIC STRUCTURE
  • Electron Configuration
    • A. Many-electron atom: An atom that contains
    • two or more electrons.
    • B. Problems with the Bohr model:
    • 1. It “assumed” quantization of the energy
    • levels in hydrogen.
    • 2. It failed to describe or predict the spectra
    • of more complicated atoms.
slide51

ATOMIC STRUCTURE

C. What are the differences in electron energy

levels in hydrogen vs. more complicated atoms?

3s 3p3d

Energy

2s

2p

Ground State Hydrogen Atom

1s

slide52

ATOMIC STRUCTURE

Splitting of the Degeneracy

2p

2s

2p

2s

Energy

1s

1s

Li

H

slide53

ATOMIC STRUCTURE

Splitting of the Degeneracy

1. In hydrogen, all subshells and orbitals in a

given principal energy level have the same energy.

They are said to be Degenerate.

2. In many-electron atoms, s-orbitals have lower

energy than p-orbitals which have lower energy

than d-orbitals which have lower energy than

f-orbitals, etc., etc.

3. Reason: Complex electrostatic interactions.

slide54

-

-

-

-

-

+

-

++

Hydrogen

+++

Helium

Lithium

A. Shielding Effect - A decrease in the nuclear force

of attraction for an electron caused by the presence

of other electrons in underlying orbitals.

B. Effective Nuclear Charge - A positive charge

that may be less than the atomic number. It is the

charge “felt” by outer electrons due to shielding by

electrons in underlying orbitals.

slide55

ATOMIC STRUCTURE

The Pauli Exclusion Principle - No two electron in

the same atom can have the same four quantum

numbers.

H + e- H -

Quantum Electron 1 Electron 2

Number

n 1 1

l 0 0

ml 0 0

ms +1/2 -1/2

slide56

The Aufbau Principle - A procedure for “building up”

the electronic configuration of many-electron atoms

wherein each electron is added consecutively to the

lowest energy orbital available, taking into account

the Pauli exclusion principle.

Order of Filling -

1s 2s 2p 3s 3p 4s 3d 4p 5s

Increasing Energy

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f 5g

mnemonic device

slide57

Designating Electron Configurations -

  • Standard Designation

H 1s1

He 1s2

Li 1s2 2s1

Be 1s2 2s2

B 1s2 2s2 2p1

C 1s2 2s2 2p2

  • Orbital Diagram Designation

H

He

Li

Be

B

C

1s

2s

2s

2p

1s

1s

1s

2s

2p

1s

2s

1s

slide58

Core Designation - A designation of electronic

configuration wherein the outer shell electrons

are shown along with the “core” configuration of

the closest previous noble gas.

Li

Na

K

Rb

[He] 2s1

[He] 2s2

Be

Mg

Ca

Sr

[Ne] 3s1

[Ne] 3s2

[Ar] 4s1

[Ar] 4s2

[Kr] 5s1

[Kr] 5s2

slide59

Hund’s Rule of Maximum Multiplicity - Electrons

occupy a given subshell singly and with parallel spins

until each orbital in the subshell has one electron.

“Electrons try to stay as far apart as possible”

  • Elevator Analogy
  • Bus Seat Analogy

[He] 2s2 2p1

[He]

B

C

N

[He] 2s2 2p2

[He]

[He] 2s2 2p3

[He]

2s

2p

slide60

Assignment: Write the electron configuration using

all three types of designation for lead (Pb).

Pb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2

4f14 5d10 6p2

Pb [Xe] 6s2 4f14 5d10 6p2

Electronic Configuration for postive ions (cations) -

Cations are formed by removing electrons in order

of decreasing n value. Electrons with the same n

value are removed in order of decreasing l value.