CHAPTER 3: ATOMIC THEORIES
3.1: Early History of Atomic Theories
-Dalton’s Atomic Theory
-smallest piece of matter is an atom; atoms are indivisible
-all atoms of an element are the same; different element have different atoms
-analogy of billiard balls
-Thomson Atomic Model
-atom composed of electrons embedded in a positively charged sphere
-analogy of the raisin bun ( plum pudding)
-Rutherford Atomic Theory – see Summary p 165
-gold foil experiment and deflection of alpha particles
-positive charge in small volume of atom; called nucleus: nucleus has most of atom’s mass
-electrons occupy most of the atom’s volume, but mass is negligible; electrons orbit nucleus
electron: ____________________________________________________________________________
proton: _____________________________________________________________________________
neutron: ____________________________________________________________________________
isotope: ____________________________________________________________________________
3.2: A Canadian nuclear physicist. Read for interest
3.3: Origins of the Quantum Theory. Read for interest. Will be covered in 12 U physics.
3.4: The Bohr Atomic Theory
- Rutherford’s model problematic because electrons would lose energy and collapse into the nucleus
- Bohr’s first postulate (see page p.176) - electrons don’t emit energy as they radiate around nucleus
- Bohr’s second postulate (see pg. 176) – electron transition – electron gains energy (in the form of a
photon) and jumps to a higher energy state, if it loses energy it jumps down to a lower energy state.
- Bohr’s evidence: Bright (emission) and dark (absorption) line spectra of the elements - only certain
photon energies can be absorbed/emitted by an atom = energy states
The Successes and Failures of the Bohr Model:
-reasonably explains periodic table
-works well for H-atom, but decreases in accuracy as you increase the atomic number
-able to predict UV and infrared patterns
-incorporated idea of quantum chemistry/physics
Quantum: ___________________________________________________________________________
valence electron : ______________________________________________________________________
Section 3.4 questions page 180 #6
3.5: Quantum Numbers – read for interest
3.6: Atomic Structure and the Periodic Table
orbital: ______________________________________________________________________________
ground state: _________________________________________________________________________
energy level diagram: __________________________________________________________________
-1s>2s>2p>3s>3p>4s>3d>4p>5s>4d>5p>6s…
Rules for filling orbitals:
-
-assume ground state
-Pauli Exclusion Principle – max of 2 electrons per orbital(=opposite spin)
-Hund’s Rule – spread electrons through similar orbital before pairing
-Aufbau Principle – fill orbital of lower energy first
eg) Aluminum =
Creating Energy Level Diagrams for Cations
-cations are missing electrons for their ground state – eg) Ca 2+ lost 2 electrons
-must remove the electrons from diagram, starting with highest quantum number - eg) 4s – 4=quantum #
-see summary p.191
Practise pg. 191 #3,4
Electron Configuration
-same info as energy level diagrams
-not pictoral, only written in a single line
eg) – aluminum = 1s22s22p63s23p1
Shorthand Form of Electron Configuration
-last noble gas + remaining energy levels, because noble gas = stable
eg) aluminum, Al = [Ne]3s23p1
Explaining the Periodic Table
-similar outer shell electron configurations apply to most families
3.7: Wave Mechanics and Orbitals
-De Broglie – particles can behave as a wave
Schrodinger – electrons behave like a wave = quantum mechanics
-Heisenburg’s Uncertainty Principle
-electron probablility density – represented by “cloud” drawings
3.8: Applications of Quantum Mechanics: Read for interest
Careers in Chemistry: p 207.
Problem Set: (see p. 188) # 8, 10, 11, 12, 13, 15