Chapter 6Electronic Structure

Section 6.5 to End

Sec 6.5 Quantum Mechanics and Atomic Orbitals

• Wave functions – describes the behavior of the electron, denoted with the Greek letter, ψ

• The wave function has a known energy, but the electron location is unknown, so the probability of its position in space is given by probability density, ψ2

• Electron density – distribution map of the probability of finding the electrons at the points of space (probability density); high probability density=high electron density

Orbitals and Quantum Numbers

• Orbitals – specific distribution of electron density in space (given by probability density); quantum mechanical model

Orbitals and Quantum Numbers

principal quantum number, n, relates to the size and energy of the orbital: integral values of 1,2,3, etc. an increase in n would mean a larger orbital, farther from the nuclear, and more energy (less tightly bound to nucleus)

There is no n= 0 value.The ground state is n = 1Excited states are n = 2, 3, 4 etc

Quantum Numbers

Azimuthal quantum number, l

* defines shape of the orbital

* designated by letters s, p, d, and f

s = 0

p = 1

d = 2

f = 3

Quantum Numbers

Magnetic quantum number, ml

* describes orientation of orbital in space* ranges from l and –l

Ex: if l = 3 then ml could be -3, -2, -1, 0, 1, 2, 3

Take a look at an orbital diagram. Compare the number of ml options with the number of

boxes per sublevel.

Quantum Numbers

Spin quantum number, ms * describes the direction of electron spin

There are 2 options for the ms value + ½ or - ½

The first electron in the orbital is spin up and the second electron is spin down.

• Electron shells – a set of orbitals with the same value of n, such as 3s, 3p, 3d

• Subshell – set or orbitals that have the same “n” and “l” values

The shell with the principal quantum number n, will have exactly n subshells

Each subshell has a specific number of orbitals. For a given l, there are 2l + 1 allowed values of m1

The total number of orbitals in a shell is n2

• Ground State – when the electron is in the lowest energy orbital

•

• Excited State – when the electron is in any other orbital

Representations of Orbitals

S- orbital

• Appears to be spherical

• Size increases as n increases

• All s-orbitals are spherically symmetrical– Nodes = the intermediate regions where ψ2

goes to zero; the number of nodes increases with increasing values of n

P - orbitals

P- orbitals

• Concentrated on two sides of the nucleus, separated by a node at the nucleus, “two lobes”

• the orbitals of a given subshell have the same size and shape but differ in spatial orientation

D - orbitals

d- and f- orbitals

• The different “d” orbitals in a given shell have different shapes and orientations in space

• When “n” is equal to or greater than 3, the d-orbitals are present. There are 5 d orbitals

• When n is equal to or greater than 4, there are 7 equal f-orbitals present

Sec 6.7 Orbitals in Many-Electron Atoms

• The presence of more than one electron greatly changes the energies of the orbitals

• The electron-electron repulsions cause different subshells to be at different energies

Effective Nuclear Charge

Screening effect the effect of inner electrons decreasing the nuclear charge experienced by outer electrons

Effective Nuclear Charge

• Each electron is simultaneously attracted to the nucleus and repelled by the other electrons

• Energy of the electron can be estimated by how it interacts with the average environment (created by the nucleus and other electrons)

Effective Nuclear Charge

Zeff = Z – S

Z is # of protons in the nucleus

S is avg # of electrons between the nucleus and the electron in question.

Energies of Orbitals

• In a many-electron atom, for a given value of n, Zeff decreases with increasing value of l

• The energy of an electron depends on the effective nuclear charge, Zeff

• In a many-electron atom, for a given value of n, the energy of an orbital increases with increasing value of l

• Degenerate - orbitals with the same energy (i.e. there are 3 p orbitals or 5 d orbitals)

Electron Spin and the Pauli Exclusion Principle

• Electron spin is a property of the electron that makes it behave as though it were a tiny magnet. The electron behaves as if it were spinning on its axis, electron spin is quantized

• Electron spin quantum number, ms , is a quantum number associated with the electron spin; two possible values + ½ or - ½

Pauli’s Exclusion Principle

• Pauli’s Exclusion Principle states that no two electrons in an atom can have the same values for n, l, ml, and ms

• Places a limit of two on number of electrons that can occupy any one atomic orbital

Sec 6.8 Electron Configurations

• Electron configuration way in which the electrons are distributed among the various orbitals of an atom

• Represent electron configuration through an orbital diagram

• each orbital represented by a box and each electron by a half arrow

Sec 6.8 Electron Configuration

Electron Configurations

• Practice writing electron configurations.

• Practice writing abbreviated electron configurations.

Sec 6.8 Electron Configuration

• Paired – electrons in the same orbital

• Unpaired – electron alone in an orbital

• Hund’s rule – for degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin in maximized

• Valence electrons - outer-shell electrons

• Core electrons- electrons in inner shells

Period 4 and Beyond

• Elements known as:* Transition Elements (and Metals)

Elements in which the “d” orbitals are filled

* Lanthanide ElementsElements in which the 4f sub shell is partly occupied

* Actinide ElementsElements in which the 5f orbitals are partly occupied

Sec 6.9 Electron Configurations and the Periodic Table

e- configuration exceptions

• If you apply the Aufbau principle to all known elements you will find that their electron configurations do not always agree with those determined experimentally. These exceptions do not have any major chemical consequences, though they are interesting to chemists.

Exceptions

Cu

The normal configuration:

1s2 2s2 2p6 3s2 3p6 4s2 3d9

The exception:

1s2 2s2 2p6 3s2 3p6 4s1 3d10

• Why do you think this happens?

Exceptions

Cr

The normal configuration:

1s2 2s2 2p6 3s2 3p6 4s2 3d4

The exception:

1s2 2s2 2p6 3s2 3p6 4s1 3d5

• Why do you think this happens?

Iron

• What are the 2 most common oxidation states of Iron?

• Write electron configurations to support the existence of the 2 most common oxidation states.

Brown & LeMay Problems

• 7-9, 13, 16, 21, 22, 27, 31, 42-46, 61, 67, 71, 74, 79, 86