Chapter 6

Electronic Structure and

Periodicity

Objectives

• Periodic law• Organization of the periodic table• Electromagnetic Radiation (added)• Quantum mechanics• Periodic Trends

Periodic Table

• Label appropriately…• Properties of elements

– Know the trends…• Explained by organization in atom

– Electronic structure

Electromagnetic Radiation

• Electromagnetic spectrum• Radio waves cosmic rays• Visible light is small portion

Radiant Energy• WAVES– Light travels in

• Light is a form of

– E.R. is a form of

Radiant Energy• All waves can be described in 4

ways:– Amplitude

•

Radiant Energy• Wavelength ( )l

• Distance between

• Measured in

Radiant Energy• Frequency (n)

– Number of

– Measured in

Radiant Energy• Speed

– Speed of light in air is – Electromagnetic radiation moves

through a vacuum at a speed

– Represented by ‘c’

As in c =

Radiant Energy• Since light moves at a constant

speed there is a relationship between wavelength and frequency

Practice

• What is the frequency of light if the wavelength is 4.65 x 10-7m?

Practice 2

• What is the wavelength of light that has a frequency of 2.88 x 1017 s-1?

Radiant Energy• Electromagnetic Spectrum

– Includes all – Visible spectrum

• • What we can see•

• Note Electromagnetic spectrum, identifying key areas (wavelength, frequency)

Electromagnetic Spectrum

Model of the Atom

• Niels Bohr• • Explains

The Bohr Atom• Model didn’t seem to work

• Did not explain

• Bohr’s Proposed Model:– The single electron of the hydrogen atom can

circle the nucleus only in

– The location of the electron(s) in each orbit determines a

– The orbit closest to

– Potential Energy increases

– Each orbit (energy level)

Another Look at the Atom

First Law of Thermodynamics

• Energy is • Energy in

• Bohr Model Animation

Another Look at the Atom

Emission Spectra

• Radiation emitted by an

• Color is • Fireworks

LineSpectrum

• Elements in gaseous states

– High temperature or high voltage

– Always

– Each element is •

Line Spectrum• Ground state

– • Excited state

–

– Farther – Higher

Line Spectrum• Electron

• Color of light

Line Spectrum• Each band of color is produced by

light of a different • Each particular wavelength has a

definite

• Each line must therefore be produced by

Beginnings…• Max Planck (1858-1947)

– Proposed that there is a fundamental restriction on the amounts of energy that an object emits or absorbs, and he called each of these pieces of energy a

• Energy is released in

Beginnings• A quantum is a finite quantity of energy

that can be gained or lost by an atom

•

Practice• Determine the energy of light that has a

frequency of 3.5 x 1016 s-1

Quantum Theory

Schroedinger, Einstein, Planck…etc…

Another Look at the Atom

• Quantum Theory describes mathematically the

• This model applies to

Terms…• Electron shell:

A region of space

There are Corresponds to the Can hold

Shells and Electrons

Shell Number (n) Electron Capacity

Terms…• Electron subshell

A region of space within a

There are the same number of

Terms• Orbital – a 3-dimensional region

– A bit more specific than • The sizes and shapes of the electron

clouds

• An individual electron cannot be pinpointed – Heisenberg’s Uncertainty Principle

Terms…• Probability and Orbitals

– Electron cloud is

– Figure 6.8 and 6.9 pg 194• Learn the rules and quantum

numbers and then put it all together…

Quantum numbers• Numbers that specify the properties of

atomic orbitals and their electrons– A system of – A method of giving – Tell you how the – Schroedinger figured this out using BIG math– If you’re curious… go HERE

• There are 4

Principal Quantum Number• The first • Energy level –

• Corresponds to the

• Value of • Symbolized by n

Azimuthal Quantum Number

• The second quantum number• Tells you the

• Usually shown as

OrbitalsEach subshell (orbital) has a specific shape

http://daugerresearch.com/orbitals/

Quantum Numbers• Orbital (azimuthal) Quantum Number:

– Indicates the – (subshell or sublevels)– Principal Quantum # Orbital Quantum #

1 2 3 4

Magnetic Quantum Number

• The third number• Tells you

• • Has values of +l to -l

S orbital

Pz orbital

Py orbital

Px orbital

D orbital

S, P, D orbitals

F orbital

Orbitron

• For a full view of the different orbital shapes, visit

• http://www.shef.ac.uk/chemistry/orbitron/index.html

Spin Quantum Number• Final quantum number• Tells you in which

• Electrons behave like little

• Value of + or – ½ or…• If you’re curious… go HERE

Putting it all together

• Table 2.6 and 2.7• Orbital arrangement on the periodic

table• Electron configuration• Aufbau principal

– “Building up’

Pauli Exclusion Principal

• No more than

• No electron can

Hund’s Rule

• Electrons will fill empty orbitals before they

• Applies to • Electrons (little magnets)

Rules for Orbital Filling• Pauli’s Exclusion Rule

– No two electrons

• Hund’s Rule

1s 2s 2p 3s 3p

Rules for Orbital Filling• Aufbau

– The order of fillingis from the

bottom (low energy) up

– Due to energy levels

Rules for Orbital Filling

Diagonal RuleThe order of filling

once the d & f sublevels are being filled

Due to energy levels

4f

Alternates…

Application of Quantum Numbers

• Several ways of writing the

• Lowest energy levels are • Electron Configuration: using the diagonal

rule, the principal quantum number (n), and the sublevel write out the location of all electrons

12C: 32S:

Definitions of Configurations• Electron Configuration

– Uses principle and orbital quantum numbers, superscript for electrons

• Orbital-filling Electron Diagram– Uses Hund’s Rule– Show electrons with

• Electron Dot Diagram– Shows ONLY the OUTER SHELL electrons

• Chemical Symbol represents the nucleus and all other electrons

• Use dots for outer shell electrons