Chapter 4.1 The Development of a New Atomic Model

Properties of Light Light as a wave: Visible light is a type of electromagnetic radiation, along with X-rays, ultraviolet and infrared light, microwaves, and radio waves. These form the electromagnetic spectrum. Waves have a repetitive nature and can be measured by wavelength () & frequency (). Wavelength unit is cm or nm. Frequency unit is waves/sec or hertz (Hz).

c = speed of light or 3.00 x 108 m/s Frequency and wavelength are related to each other through the following equation: c =  c = speed of light or 3.00 x 108 m/s As wavelength increases, frequency decreases and vice versa.

Change cm to m and plug into equation Example Determine the frequency () of light whose wavelength () is 6.87 x 10-8 cm. c =  c = 3.00 x 108 m/s = c/ Change cm to m and plug into equation = 3.00 x 108 m/s 6.87 x 10-10 m  = .437 x 1018 ≈ 4.37 x 1017 Hz

Light as a particle: Photoelectric effect is the emission of electrons from metal when light shines on it. A quantum of energy is the minimum amount of energy that can be lost or gained by an atom. A photon is a particle of electromagnetic radiation with no mass and carrying a quantum of energy.

Ch 4.2 Notes The Quantum Model of the Atom Objectives: To describe the quantum mechanical model of the atom. To describe the relative sizes and shapes of s and p orbitals.

Electrons as Waves Behavior of electrons is similar to the behavior of waves. Electron waves can only exist at certain frequencies.

Heisenberg Uncertainty Principle Werner Heisenberg had an idea on how to detect the location of electrons. Heisenberg Uncertainty Principle: it is impossible to determine simultaneously both the position and velocity (speed) of an electron.

The Schrödinger Wave Equation Developed an equation that treated electrons in atoms as waves and quantization of electron energies was an outcome of the equation. Quantum Theory: mathematically describes the wave properties of electrons and other very small particles.

Principal Quantum Number ( n ) Main energy level (shell) Size of the orbital PERIOD # Number of orbitals per main energy level is equal to n2. Number of electrons = 2n2. 1s 2s s Orbitals – Orbitals with l = 0 are s orbitals and are spherically symmetrical, with the greatest probability of finding the electron occurring at the nucleus. – All orbitals with values of n > 1 and l  0 contain one or more nodes. – Three things happen to s orbitals as n increases: 1. they become larger, extending farther from the nucleus 2. they contain more nodes 3. for a given atom, the s orbitals become higher in energy as n increases due to the increased distance from the nucleus 3s Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Angular Momentum Quantum Number Indicates the shape of the orbital, represented by l. Values of l allowed are 0 and all positive integers less than or equal to n -1. Ex. If n=2, shapes are l = 0 and l = 1 A sublevel consists of orbitals in a main energy level with the same value of l.

Magnetic Quantum Number Indicates the orientation of an orbital around the nucleus, represented by m. Values of m are whole numbers, including zero, from –l to +l

Electron capacities Electron capacities Copyright © 2006 Pearson Benjamin Cummings. All rights reserved.

Spin Quantum Number A single orbital has a maximum of two electrons, and those electrons must have opposite spins. It has two values +1/2 and -1/2

Ch 4.3 Electron Configurations Electron Configurations: the arrangement of electrons in an atom. Each element has a unique electron configuration. Ground-state Electron Configuration is the lowest energy arrangement of the electrons for an element.

General Rules Aufbau Principle Electrons fill the lowest energy orbitals first. “Lazy Tenant Rule”

General Rules WRONG RIGHT Hund’s Rule Within a sublevel, place one electron per orbital before pairing them. “Empty Bus Seat Rule” WRONG RIGHT Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

General Rules Pauli Exclusion Principle Wolfgang Pauli Pauli Exclusion Principle Each orbital can hold TWO electrons with opposite spins. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Electron Filling in Periodic Table s s p 1 2 d 3 4 5 6 7 f

Elements in the s - blocks really should include He, but He has the properties of the noble gases, and has a full outer level of electrons. 31

The P-block p1 p2 p3 p4 p6 p5 32

Transition Metals - d block Note the change in configuration. s1 d5 s1 d10 d1 d2 d3 d5 d6 d7 d8 d10 33

F - block f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14 f13 Called the “inner transition elements” f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14 f13 34

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1s1 Periodic Patterns 1st Period s-block Example - Hydrogen # element in block 1st Period s-block Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Periodic Patterns p s d (n-1) f (n-2) Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

1s2 2s2 2p3 N Notation 1s 2s 2p 7e- N Orbital Diagram 7 14.0067 7 Notation Orbital Diagram 1s 2s 2p N 7e- Electron Configuration 1s2 2s2 2p3 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

Electron Filling in Periodic Table s s p 1 2 d 3 K 4s1 Ca 4s2 Sc 3d1 Ti 3d2 V 3d3 Cr 3d4 Cu 3d9 Cr 3d5 Mn 3d5 Fe 3d6 Co 3d7 Ni 3d8 Cu 3d10 Zn 3d10 Ga 4p1 Ge 4p2 As 4p3 Se 4p4 Br 4p5 Kr 4p6 4 Cr 4s13d5 Cu 4s13d10 Cr 4s13d5 4s 3d Cu 4s13d10 4s 3d

Order in which subshells are filled with electrons (Fig 19 pg 116) 2 2 6 2 6 2 10 6 2 10 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d …

Shorthand Electron Configuration Ge 72.61 32 Shorthand Electron Configuration Example - Germanium [Ar] 4s2 3d10 4p2 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

S 16e- 1s2 2s2 2p6 3s2 3p4 S 16e- [Ne] 3s2 3p4 Notation Core Electrons 32.066 16 Notation Longhand Configuration S 16e- 1s2 2s2 2p6 3s2 3p4 Core Electrons Valence Electrons Shorthand Configuration S 16e- [Ne] 3s2 3p4 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem