Models of the Atom/ Quantum Theory

1. Models of the Atom/ Quantum Theory John Dalton: believed the atom was a solid mass J.J Thomson: Plum Pudding Model In 1903, J. J. Thomson proposed a subatomic model of the atom. The model pictured a positively-charged atom containing negatively-charged electrons. Thomson visualized electrons in homogeneous spheres of positive charge in a way that was analogous to raisins in English plum pudding. Thus, the Thomson proposal became popularly know as the plum pudding model of the atom.

3. Ernest Rutherford: In 1911, Rutherford proposed a new model of the atom. He suggested that negatively-charged electrons were distributed about a positively-charged nucleus. Most of the mass of the atom was concentrated in the nucleus, the rest was empty space.

4. Niels Bohr: In 1913 Bohr produced a model of the atom in which radiation was emitted only when an electron jumped from one orbit, or shell, to another. The frequency of light that an atom emitted was not related to the frequency of or in the atom, rather, it was based on the difference between two energy levels within the atom. Electrons can jump from a low-energy orbit near the nucleus to orbits of higher energy by absorbing energy (green trails). When the electrons return to a lower energy level (purple trails), they release the excess energy in the form of radiation of a characteristic wavelength, such as visible light. *no e- between energy levels *all e- have fixed energy to stay in their energy level *levels are not evenly spaced

6. Louis de Broglie was born in 1892. He is best known for his theory, which he confirmed in 1927, that matter has the properties of both particles and waves. This theory was later used by Erwin Schrödinger to develop wave mechanics. When Schrödinger studied the atom, he abandoned the idea of precise orbits, and replaced them with a description of the regions of space, which he called orbitals, where the electrons were most likely to be found. This is the atomic model of the atom that we use today.

7. Quantum Numbers: Principle Quantum Number(n) -energy level n=1, 2, 3, etc. -the distance from the nucleus increases as n increases -the greatest number of electrons possible in any one energy level is 2n2. Ex. n=2 8 electrons n= 4 32 electrons -n numbers match up to period numbers

8. Second Quantum Number (l) -represents the sublevels with an energy level, shapes # of sublevels = principle quantum level l = 0 to (n –1) If l equals 0 = s shape 1 orbital, 2 e- 1 = p shape 3 orbitals, 6 e- 2 = d shape 5 orbitals, 10 e- 3 = f shape 7 orbitals, 14 e-

11. Energy Levels and Orbitals diagram Nucleus

12. nlshapes on that energy level#e-

13. Third Quantum Number, ml ml = +l to –l -represents atomic orbital: space occupied by one pair of electrons Fourth Quantum Number, ms ms = + ½ or – ½ -represents the spin of the e- -if 2 electrons occupy the same orbital they must have opposite spins

14. Write all of the quantum numbers associated with the electrons in Mg.

15. Electron Configuration: ways in which electrons are arranged around the nuclei of atoms • 3 rules of configuration: • Aufbau Principle: Electrons enter orbitals of lowest energy first, • arrow diagram • Pauli Exclusion Principle: An atomic orbital may describe atmost two electrons. No two electrons will have all four identical quantum numbers. (opposite spins) • Hund’s Rule: When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel spins.

16. Exceptions: Cr, Cu Mo, Ag

17. Electron Dot Diagrams -include s and p orbitals from the highest energy level Valence Electrons: electrons involved in bonding.