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Unit 5- Electrons in Atoms Quantum Model of the Atom Reminder Thomson Model - Plum Pudding Rutherford Model - Nuclear model 1912-1913 Rutherford gathered physicists, including Neils Bohr to work on WHY atoms don’t collapse (negative eattract positive p+) Theory Planks Quantum Theory suggested the eexist in energy levels that have discrete amounts of energy. e- can jump(quantum leap) to higher levels after gaining a Quantum of energy. Same amount of energy is given off upon return to GROUND STATE. Energy Levels. A ladder represents the idea. Lowest rung represents energy level 1, n= 1 (closest to the nucleus) Each successive rung is a higher energy level. ( n=2, n=3, n=4…) The energy level is the most likely location an electron can be found within the cloud. Energy Levels n = 1 can hold 2 e- max n = 2 can hold 8 e- max n = 3 can hold 18 e- max n = 4 can hold 32 e- max Try these: Mg atomic number 12 (12 p+ and 12 e-) n = 1 2 e- n = 2 8 e- n = 3 2 e- Valence electrons The outer most electrons are called VALENCE ELECTRONS They are the bonding electrons – VERY IMPORTANT B. Quantum Mechanics Schrödinger Wave Equation (1926) finite # of solutions quantized energy levels defines probability of finding an e- Ø 1s = () e 1 Z 3/2 -ó ∂ a0 B. Quantum Mechanics Orbital (“electron cloud”) Region in space where there is 90% probability of finding an e- Orbital Radial Distribution Curve C. Quantum Numbers- Address of an e- Four Quantum Numbers: Specify the “address” of each electron in an atom UPPER LEVEL C. Quantum Numbers 1. Principal Quantum Number ( n ) Energy level Size of the orbital n2 = # of orbitals in the energy level C. Quantum Numbers 2. Angular Momentum Quantum # ( l ) Energy sublevel Shape of the orbital s p d f C. Quantum Numbers n = # of sublevels per level n2 = # of orbitals per level Sublevel sets: 1 s, 3 p, 5 d, 7 f C. Quantum Numbers 3. Magnetic Quantum Number ( ml ) Orientation of orbital Specifies the exact orbital within each sublevel C. Quantum Numbers px py pz C. Quantum Numbers 4. Spin Quantum Number ( ms ) Electron spin +½ or -½ An orbital can hold 2 electrons that spin in opposite directions. B. Notation Orbital Diagram (drawing of e- location) O 8e- 1s 2s 2p Electron Configuration (distribution of e- in the orbitals) 2 2 4 1s 2s 2p C. Orbital Shapes S shape orbital C. Orbital Shapes p shaped orbitals (dumbbells) px py pz C. Orbital shapes d shaped orbitals (cloverleaf) E- configuration Electrons fill the atom from Low energy to High energy. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6 How do we keep it straight? Feeling overwhelmed? Read B. Notation Longhand Configuration P 15e- 1s2 2s2 2p6 3s2 3p3 Core Electrons Valence Electrons Shorthand Configuration P 15e 2 3 [Ne] 3s 3p B. Notation Longhand Configuration Ca 20e- 1s2 2s2 2p6 3s2 3p6 4s2 Core Electrons Valence Electrons Shorthand Configuration Ca 20e 2 [Ar] 4s C. Periodic Patterns s p 1 2 3 4 5 6 7 f (n-2) d (n-1) 6 7 © 1998 by Harcourt Brace & Company A. General Rules Aufbau Principle Electrons fill the lowest energy orbitals first. A. General Rules Pauli Exclusion Principle Each orbital can hold TWO electrons with opposite spins. A. General Rules Hund’s Rule Within a sublevel, place one e- per orbital before pairing them. WRONG RIGHT