Chapter 4 Arrangement of Electrons in Atoms

I. The Development of a New Atomic Model H Electromagnetic Radiation: H Electromagnetic Spectrum: H Electromagnetic Radiation: H Electromagnetic Spectrum:

H Wavelength ( ): corresponding points on adjacent waves---Ex: Frequency ( ): # of waves that pass a point in a specific time H c = ( ) ( ) inversely proportional H Wavelength ( ): corresponding points on adjacent waves---Ex: Frequency ( ): # of waves that pass a point in a specific time H c = ( ) ( ) inversely proportional

c : m/s : m, cm, nm : waves/second--Hertz (Hz) H c = ( ) ( ) inversely proportional c : m/s : m, cm, nm : waves/second--Hertz (Hz)

H Photoelectric Effect: emission of e- by certain metals when light shines on them

H Quantum: min quantity of nrg that can be lost or gained by an atom H E = (h) ( ) o J = (J  s) (Hz) o Planck’s constant: X J  s H Quantum: min quantity of nrg that can be lost or gained by an atom H E = (h) ( ) o J = (J  s) (Hz) o Planck’s constant: X J  s

Einstein o dual wave-particle to describe light H Photon: radiation with zero mass carrying a quantum of nrg o packet of nrg emitted when an e- drops nrg levels Einstein o dual wave-particle to describe light H Photon: radiation with zero mass carrying a quantum of nrg o packet of nrg emitted when an e- drops nrg levels

H Ground state: lowest nrg state H Excited state: higher potential nrg H Ground state: lowest nrg state H Excited state: higher potential nrg

--Bohr’s Model-- G e - exist only in orbits with specific amounts of energy called energy levels G Therefore… G e - can only gain or lose certain amounts of energy G only certain photons are produced G e - exist only in orbits with specific amounts of energy called energy levels G Therefore… G e - can only gain or lose certain amounts of energy G only certain photons are produced

Line-Emission Spectrum ground state excited state ENERGY IN PHOTON OUT

Bohr Model Energy of photon depends on the difference in energy levels -Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom

Other Elements G Each element has a unique bright-line emission spectrum. G “Atomic Fingerprint” Helium zBohr’s calculations only worked for hydrogen!  ----pg 97

II. The Quantum Model of the Atom G A. Electrons as Waves o Diffraction: bending of a wave as it passes by the edge of an object o Interference: results when waves overlap G A. Electrons as Waves o Diffraction: bending of a wave as it passes by the edge of an object o Interference: results when waves overlap

EVIDENCE: DIFFRACTION PATTERNS VISIBLE LIGHT ELECTRONS

zHeisenberg Uncertainty Principle yImpossible to know both the velocity and position of an electron at the same time

G Schrödinger Wave Equation (1926) G finite # of solutions  quantized energy levels G defines probability of finding an e - G Schrödinger Wave Equation (1926) G finite # of solutions  quantized energy levels G defines probability of finding an e -

A. Atomic Orbitals and Quantum Numbers G Orbital: probable location of an e- G Quantum #: properties of atomic orbitals and properties of e-’s in orbitals G Principal quantum #: (n), indicates main nrg level occupied by the e- o n = occupies 1st nrg level G Orbital: probable location of an e- G Quantum #: properties of atomic orbitals and properties of e-’s in orbitals G Principal quantum #: (n), indicates main nrg level occupied by the e- o n = occupies 1st nrg level

G Angular momentum quantum #: (l), indicates shape of orbital G Magnetic quantum #: (m), orientation of an orbital G Spin quantum #: which spin state (+)(-) G ***See table 4-2 pg 104 G Angular momentum quantum #: (l), indicates shape of orbital G Magnetic quantum #: (m), orientation of an orbital G Spin quantum #: which spin state (+)(-) G ***See table 4-2 pg 104

zOrbital (“electron cloud”) yRegion in space where there is 90% probability of finding an e - Orbital Radial Distribution Curve

zFour Quantum Numbers: ySpecify the “address” of each electron in an atom UPPER LEVEL

1. Principal Quantum Number ( n ) yEnergy level ySize of the orbital yn 2 = # of orbitals in the energy level

2. Angular Momentum Quantum # ( l ) yEnergy sublevel yShape of the orbital s p d f

zn=# of sublevels per level zn 2 =# of orbitals per level zSublevel sets: 1 s, 3 p, 5 d, 7 f

3. Magnetic Quantum Number ( m l ) yOrientation of orbital  Specifies the exact orbital within each sublevel

4. Spin Quantum Number ( m s ) yElectron spin  +½ or -½ yAn orbital can hold 2 electrons that spin in opposite directions.

III. Electron Configuration G Aufbau principle: lowest nrg orbits fill first G Pauli exclusion principle: no 2 e-’s can have the same 4 quantum #’s. This is where spin allows 2 e-’s to be in the same orbit o Ex:  G Aufbau principle: lowest nrg orbits fill first G Pauli exclusion principle: no 2 e-’s can have the same 4 quantum #’s. This is where spin allows 2 e-’s to be in the same orbit o Ex: 

G Hund’s rule: orbital of equal nrg are occupied by 1 e-, before any is occupied by 2 e-’s o Ex:    G Orbital Notation: ex: pg 107 G Electron Config Notation: pg 107 G Electron Dot diagram: ex G Hund’s rule: orbital of equal nrg are occupied by 1 e-, before any is occupied by 2 e-’s o Ex:    G Orbital Notation: ex: pg 107 G Electron Config Notation: pg 107 G Electron Dot diagram: ex

G Noble gases: G are inert G complete octet G --show ex---- G Noble gases: G are inert G complete octet G --show ex----

G Table 4-3 pg Principal #  energy level 2. Ang. Mom. #  sublevel (s,p,d,f) 3. Magnetic #  orbital 4. Spin #  electron G Table 4-3 pg Principal #  energy level 2. Ang. Mom. #  sublevel (s,p,d,f) 3. Magnetic #  orbital 4. Spin #  electron

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