Chapter 6

Modern Atomic Theory

Review…

• Dalton• Thomson• Rutherford

– Model doesn’t explain how the negative electron can stay in orbit and not be attracted to the positive proton

Electromagnetic Radiation

• Light travels in

• Light is a form of – Form of energy that exhibits

Electromagnetic Radiation

• All waves can be described in 3 ways:– Amplitude –

– Wavelength ( ):l

– Frequency ( ):n

Electromagnetic Radiation

• Speed of light in air: Electromagnetic radiation moves through a vacuum at speed of

• Since light moves at constant speed there is a relationship between wavelength and frequency:

Wavelength and frequency are

inversely proportional

Electromagnetic Spectrum

Photoelectric Effect

• The emission of

– Albert Einstein (1905) used Planck’s equation to explain this phenomenon;• proposed that light consists of

• Photon =

Photoelectric Effect

• He (Einstein) explained that the photoelectric effect would not occur if the frequency and therefore

• Analogy:– 70 cents placed in soda machine: no

soda– 30 cents more and you will get your

soda

Niels Henrik David Bohr

• 1885-1962• Physicist• Worked with

Rutherford– 1912

• Studying line spectra– of hydrogen

Niels Henrik David Bohr

• 1913 – proposed new atomic structure– Electrons exist in

– Electrons

The Bohr Atom• Nucleus with • Electrons move in

• When an electron moves from one state to another the energy lost or gained is in

• Each line in a spectrum is produced when an electron moves from

The Bohr Atom

• Model didn’t seem to work with atoms with more than one electron

• Did not explain chemical behavior of the atoms

Now…

• Light can be described as

•

• What does this mean for the atom???

LineSpectrum• Elements

in gaseous states give

off colored light

– High temperature or high voltage– Always the same– Each element is unique

• Spectra

Line Spectrum• Ground state

– • Excited state

–

– –

Line Spectrum• Electron

• Color of light emitted depends on

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 emission of photons with

Line Spectrum• Whenever an excited electron

• The energy of this photon is equal to the difference

Wave Matters…

•Louis de Broglie (1924)

•Proposed that electrons might have a

•Used observations of normal wave activity

Problems…

• Wave theory does not explain– Heated iron gives off heat

• 1st red glow yellow glow white glow

– How elements such as barium and strontium give rise to green and red colors when heated

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

• Energy is released in

Beginnings

• A quantum is a finite quantity of energy that can be gained or lost by an atom

• This constant, h, is the same for all electromagnetic radiation

Bohr’s Equation

•En = (-RH)(1/n2)

• Where RH = 2.18 x 10-18J• And n = principal quantum

number, 1 to infinity

Jumping electrons…

• If an electron moves from one energy level to another, the change in energy can be determined by the following equation:

• E = Ef – Ei = hν• Or simply: E = hv

– Where h=6.626 x 10-34 J s

Then… by substitution…

ν =E

h=

RH

h( 1 1-

ni2 nf

2(

Finally… Matter waves

• All moving particles

• Some is apparent, some not.• De Broglie’s equation

λ =h

mν

Smart guy…

•Erwin Schrodinger (1926)

•Used mathematical understanding of wave behavior – devised an equation that treated electrons moving around nuclei as waves

•Quantum Theory

Uncertainty principle

• Heisenberg:

Quantum Theory

• Describes mathematically the wave properties of electrons and other very small particles

• Applies to all elements (not just H)

Quantum Numbers

• Numbers that specify the

• Principle Quantum Number:– Symbolized by n,

Energy Levels of Electrons

• Principle energy levels– Designated by letter n– Corresponds to the– Each level divided into sublevels

• 1st energy level has• 2nd energy level has• Etc.

Orbitals• Electrons don’t

• Orbital: region in space where

– Each orbital sublevel can hold

OrbitalsEach sublevel (orbital) has a specific shape

http://daugerresearch.com/orbitals/

Quantum Numbers• Orbital Quantum Number:

– Indicates the shape of an orbital– (subshell or sublevels)– s, p, d, f Principal Quantum # Orbital Quantum

#1 2 3 4

Quantum Numbers

• Magnetic Quantum Number:– Indicates the

– Orbital position with respect to

Orbitron

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

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

Orbitals• Pauli exclusion principle:

• Electrons can only spin• Shown with

Rules for Orbital Filling• Pauli’s Exclusion Rule

– No two electrons have

• Hund’s Rule– Electrons will remain

1s 2s 2p 3s 3p

Rules for Orbital Filling

• Diagonal Rule– The order of filling

once the d & f sublevels are being filled

– Due to energy levels

Rules for Orbital Filling

Application of Quantum Numbers

• Several ways of writing the address or location of an electron

• Lowest energy levels are filled first• Electron Configuration:

12C: 32S:

Application of Quantum Numbers

• Orbital filling electron diagram: using Hund’s rule and the diagonal rule write out the location of all electrons

• See examples on whiteboard