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 C = speed of light L = wavelength V = frequency

Electromagnetic Spectrum Visible spectrum – white light (what we can see) Wavelengths of visible spectrum are 350 nm to 700 nm

Photoelectric Effect The emission of Photon = 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 Safeco field example

Now… Light can be described as What does this mean for the atom???

Line Spectrum Elements in gaseous states give off colored light High temperature or high voltage Always the same Each element is unique Spectra Also known as the “atomic emission spectrum of an element” Visible light or electromagnetic radiation is emitted when an atom passes from a state of higher potential energy to a state of lower potential energy

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 Safeco field example

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

Energy is released in 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

En = (-RH)(1/n2) Bohr’s Equation 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… ( E RH ( 1 1 - = ν = ni2 nf2 h h

h λ = mν Finally… Matter waves All moving particles Some is apparent, some not. De Broglie’s equation h λ = 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 Safeco field example

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 Safeco field example

Orbitals Each sublevel (orbital) has a specific shape Safeco field example 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 Safeco field example

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 Safeco field example

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