Atomic Structure and Periodicity

Chapter 7

Section 7.1 Electromagnetic Radiation

• The electromagnetic spectrum organizes waves by wavelength, frequency, and energy

– wavelength- the length from a point on a wave to a corresponding point later in the wave.

– frequency- the number of times a full wave cycle passes by a reference point in one second.

The electromagnetic spectrum

Relationships between , , and E• Wavelength and frequency are inversely

proportional (a long wave will take longer to pass by a reference point, thus making its frequency lower).

• Energy is directly related to frequency. Higher frequency waves have more energy than lower frequency waves.

• The speed of a wave is constant in a vacuum. (3.0 x 108 m/s = c, speed of light)

* = c

Show Me Problem

• A red light emits light of about 650 nm wavelength. What is the frequency of the red light?

650 nm 6.5 x 10-7m(6.5x10-7)() = (3.0 x 108)

= 4.61 x 1014 Hz

* = c

Section 7.2 The Nature Of Matter

• Max Plank– Quantizes energy

• Matter can absorb energy, but only in whole number ratios of the term h.

Plank’s constant6.626 x 10-34 J*s

EQUATION: E = (h*)

Einstein’s Contribution

• Albert Einstein– Quantizes Radiation

• E = mc2 … Energy has mass and velocity. Electromagnetic Radiation must be made up of particles called photons.

Duality of Light: Electromagnetic radiation has the capacity to behave both as

a wave and a particle.

DeBroglie’s Contribution• Louis DeBroglie

– Determines the Duality of Matter• If waves act as particles, do particles act as waves?• Set the Einstein and the Plank equations equal to each

other.mc2 = E = hc/

m = h/ch/m

Duality of Matter: Electrons have the capacity to behave both as a particle and

a wave.

How can we be certain?• X Ray Crystallography

X Ray Diffraction Pattern of 2-terphenyl-4-yl-5-phenyl thiophene (PPPTP)

X-Ray Diffraction Pattern of Beryl.

Different color and shading patters appear

as the electron “waves” cause

diffraction: constructive and

destructive interference with the x-rays that are exposed

to the crystal.

Atomic Spectrum of Hydrogen

• Extensively studied by atomic theorists such as Bohr.• High energy sparks cause hydrogen gas molecules (H-H)

to break apart suddenly, with some electrons in higher energy levels than would be expected normally.

• As the electrons fall back to their ground states, energy is released.

• Each color in the spectrum relates to an electron in a different energy level.

• Plank’s equation can be used to determine the color of the light produced or the energy of the electron that is being observed.

More on the atomic spectrum of hydrogen

Section 7.4 The Bohr Model

• Based upon the study of the Hydrogen Spectrum, Bohr designs paths for electrons to travel while orbiting the nucleus. [ORBITS]

• Each orbit corresponded to a different energy level.

221810178.2nZxE

energy of e-energy level nuclear

charge (protons)

Dr Quantum Video

Show Me

• An electron in a hydrogen atom moves from energy level one to energy level 2. What is the change in energy the electron experiences?

E1 = -2.178 x 10-18 (12/12)= -2.178 x 10-18

E2 = -2.178 x 10-18 (12/22)= -5.445 x 10-19

E = E2-E1 = -5.445 x 10-19-(-2.178 x 10-18)

E = 1.634 x 10-18 JoulesEndo or Exo?

Does this make sense?

Section 7.5 The Quantum Mechanical Model of the Atom

• Determine that if the electron acts as a standing wave (a wave that is fixed in place), then there are only certain orientations for it to exist without causing destructive interference with itself.

Heisenberg

De Broglie Scrhödinger

Heisenberg Uncertainty Principle

• There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle.

• The more certain you are of the location of an electron, the less certain you can be of its momentum

• The more certain you are of the momentum of an electron, the less certain you can be of its position.

x * m = h/4

The Wave Equation

• Schrodinger’s wave equation is used to define the location of electrons as waves.

• A wave function is called an orbital. – ORBITALS = ORBITS

• Wave functions are impossible to visualize. We picture the electron density map (aka: electron probability diagram)

H = E

Section 7.6: Quantum Numbers

• Quantum numbers describe the properties of orbitals.

symbol

name Values meaning

n Principal quantum number

1, 2, etc Energy level

l Angular momentum

0 to (n-1) Orbital shape

ml Magnetic quantum number

l to -l Orientationof orbital

n l orbital ml # of orbitals

1 0 1s 0 1

2 0 2s 0 1

2 1 2p -1, 0, 1 3

3 0 3s 0 1

3 1 3p -1, 0, 1 3

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

4 0 4s 0 1

4 1 4p -1, 0, 1 3

4 2 4d -2, -1, 0, 1, 2 5

4 3 4f -3, -2, -1, 0, 1, 2, 3

7

Common Orbital Shapes

Section 7.7 Orbital Shapes

• Areas of high probability are separate by areas of low probability. (NODES)

• Degenerate orbitals have different orientation or shape but the same ENERGY.

• Lowest available energy level for an electron = ground state

• Higher energy levels than expected = excited states

Section 7.8 Electron Spin and Pauli Principle

• Electrons exhibit a fourth quantum number.

• Electron spin quantum number (ms)

• Values of ½ and -½. Indicates magnetic moment of electron. Electrons can only spin in one of 2 opposite directions.

Pauli exclusion Principle:In a given atom, no two electrons

Can have the same set of Quantum numbers.

Electron Configurations

“Diagonal Rule”

5s 5p 5d 5f4s 4p 4d 4f3s 3p 3d2s 2p1s

3 Ways for Electron Configurations

Electron ConfigurationDiagrams

Long-HandConfigurations

Nobel Gas Configurations

Electron Configuration Rule Summary

• Electrons enter lowest energy orbitals first.

• Only two electrons per degenerate orbital.

• Electrons spread out among degenerate orbitals before pairing.

Exceptions to the Configuration Rules

• A fully filled orbital is more stable than a partially filled orbital.

• half-filled orbital is more stable than a more/less partially filled orbital.

Mo AgEu AmCr Cu

Copper and Chromium

Cu expected configuration:

1s22s22p63s23p64s23d9

Cu actual configuration: 1s22s22p63s23p64s13d10

Cr expected configuration:

1s22s22p63s23p64s23d4

Cr actual configuration:

1s22s22p63s23p64s13d5

Molybdenum and Silver

Mo expected configuration:

1s22s22p63s23p64s23d104p65s24d4

Mo actual configuration:

1s22s22p63s23p64s23d104p65s14d5

Ag expected configuration:

1s22s22p63s23p64s23d104p65s24d9

Ag actual configuration:

1s22s22p63s23p64s23d104p65s14d10

Europium and Americium

Eu expected configuration:

1s22s22p63s23p64s23d104p65s2

4d105p66s24f6

Eu actual configuration:

1s22s22p63s23p64s23d104p65s2

4d105p66s14f7

Am expected configuration:

1s22s22p63s23p64s23d104p65s2

4d105p66s24f145d107s25f6

Am actual configuration:

1s22s22p63s23p64s23d104p65s2

4d105p66s24f145d107s15f7

Electron Configuration Diagrams

Noble Gas Notation

Expanded Titanium : 1s22s22p63s23p64s23d2

Noble Gas Titanium: [Ar] 4s23d2

Use the noble gas that comes before the element as a benchmark, then tack on

the extra occupied orbitals.

Periodic Trends: Electron Configurations

• Counting down tells what energy orbital.• Counting over tells how many electrons.

Periodic Trends Activity

• http://academic.pgcc.edu/~ssinex/excelets/PT_interactive.xls

Go to this excel sheet and click on the bottom tab labeled “atom properties”

Periodic Trends: Atomic SizeIncreasing Atomic Size

Periodic Trends: Ionization Energy

Increasing 1st Ionization Energy

Periodic Trends: Electron AffinityIncreasing electron affinity

Ion Size

• Negative ions indicate a gain of electrons.

• They are larger than the atom from whence they are formed.

• Positive ions indicate a loss of electrons.

• They are smaller than the atom from whence they are formed.