Chapter 4.1 Radiant Energy Wave-Particle Nature of Light Electrons and light have a dual wave-particle nature. Electromagnetic Radiation (EMR) Form of energy that exhibits wavelike behavior and travels at the speed of light. Speed of Light (C) = 3 x 10 8 m/s

Components of a Wave Wavelength () lambda Units: any unit of length (m) Distance between corresponding points of a wave. Crest to Crest or Trough to Trough

Components of a Wave Frequency () nu Units: Hertz (Hz) or 1/s How often a wavelength passes a given point in time.

Components of a Wave Amplitude Height of the wavelength. Measured from the origin to crest or origin to trough. Brightness of light.

Wavelength vs. Frequency Inversely proportional. As wavelength increases, frequency decreases.

Chapter 4.1 Radiant Energy Spectrums Range of wavelengths for a series of waves. Electromagnetic Spectrum Consist of all electromagnetic radiation. Continuous Spectrum Spectrum where all wavelengths within a given range are together. Examples: Visible Light, X-Rays, U.V. Light, etc

EMR Spectrum 7 Parts Longest wavelength to Shortest: Radio Microwaves Infrared Visible Light U.V. Light X-Rays Gamma-Rays

Chapter 4.1 Radiant Energy Problems What is the wavelength of EMR that has a frequency of 7.50 x 10 12Hz?

Chapter 4.1 Radiant Energy Problems: 1. Determine the frequency of light with a wavelength of 4.257 x 10-7 cm. 2. What is the wavelength of U.V light that has a frequency of 4.50 x 10 16 Hz? 3. What is the wavelength and color of light, that has a frequency of 6.00 x 10 14 KHz?

Chapter 4.2 Quantum Theory Photoelectric Effect Emission of electrons by certain metals when sufficient light shines on them.

Chapter 4.2 Quantum Theory Photoelectric Effect

Chapter 4.2 Quantum Theory Photoelectric Effect

Chapter 4.2 Quantum Theory Finite quantity of energy that can be gained or lost by an atom. Planck’s Equation: h = 6.63 x 10 –34 Js E = quantum of energy Photon An individual quantum of light, caused by electrons losing quanta of energy. E = h 

Chapter 4.2 Quantum Theory Visible Light Emissions As electrons gain quanta of energy they release it in the form of photons. Energy States of an Atom Ground State- an atoms lowest energy level. Excited State- an atoms highest energy level. , is produced when electrons drop from the excited to the ground states. Line Spectrum

Chapter 4.2 Quantum Theory

Chapter 4.2 Quantum Theory Problems: What is the energy of U.V. light with a frequency of 4.50 x 10 16 Hz?

Chapter 4.2 Quantum Theory Problems: Determine the energy of light that has a wavelength of 450nm.

Chapter 4.2 Quantum Theory Equations:

Chapter 4.2 Quantum Theory Problems: 1. What is the energy of a photon of green light with a frequency of 5.80 x 1014 1/s? 2. What is the energy, in joules, of a quantum of radiant energy whose wavelength is 6.82 x 10 –6 cm? 3. Determine the wavelength of a photon that has 3.11 x 10 –19 J of energy. 4. Determine the frequency, in MHz, of a photon that has wavelength of 1.36 x 10 10 nm.

Summary – restricted the amount of energy that an object emits or absorbs as a quantum. – used Planck’s theory and explained the photoelectric effect. – light travels as tiny particles, photons. Planck Einstein Compton

Another Look at the Atom Chapter 4-3 Another Look at the Atom

Bohr’s Model The Line Spectra demonstrates that the energy levels of an electron in an atom are quantized Similar to the rungs of a ladder, nothing exist in between. (For Hydrogen (1 p+ & 1 e- ) 1st Energy Level n = 1 2nd & so on n = 2,3,4,5,6, … ∞ Only electrons dropping from a Higher Level to a Lower one emit EMR A Number of Possibilities for electron drops

Hydrogen’s Line Spectrum Several Series of lines are observed Electron Drops to the n = 1 Level Lyman Series (U.V. Range) Electron Drops to the n = 2 Level Balmer Series (Visible Range) Electron Drops to the n = 3 Level Paschen Series (Infrared Range)

The Lines become more closely spaced as the levels increase The Bohr model explained spectral lines but not how atoms bonded. Ultimately Displaced

1924 Louis de Brogile – French Graduate Student (asked an important question) If light behaves as waves & particles, can particles of matter behave as waves? Derived an Equation Predicts that all matter exhibits wavelike motions. h – Plank’s Con. m – mass v - velocity

Small Objects – Large Wavelengths Large Objects – Small Wavelengths 200 g Baseball @ 30 m/s -  = 10-32 cm Undetectable Small Objects – Large Wavelengths 9.11 x 10-28 g @ 30 m/s -  = 10-3 cm Very Detectable w/ proper instruments New Ballgame – Classical Mechanics vs. Quantum Mechanics New method for describing the motions of subatomic particles

STOP!

Heisenberg’s Uncertainty Principle It is impossible to know exactly both the velocity & the position of a particle at the same time. Accuracy of V  then Position 

Classical Vs Quantum Classical adequately describes the motions of bodies much larger than the atoms of which they are composed. It appears that such a body loses energy in any amount Quantum describes the motions of subatomic particles and atoms as waves. These particles gain or lose energy in packages called quanta.

Quantum Mechanical Model Modern description of the electrons derived from the mathematical solution to the Schrodinger equation. Erwin Schrodinger - used wave mechanics to show the electrons about the nucleus emit vibration frequencies that were constant. Quantum Numbers - specify the properties of atomic orbitals and their electrons. distance from the nucleus.

Chapter 4.4 Quantum Numbers Principal Quantum Number (n) Main energy level surrounding the nucleus. Size of each orbital. Primary distance from the nucleus. Has values of n =1 to 7, 1 is the closest 7 is the farthest from the nucleus.

Chapter 4.4 Quantum Numbers Orbital Quantum Number (l) Shape of the orbitals. Referred to as subshells.

Chapter 4.4 Quantum Numbers s orbital p orbital d orbital f orbital

Chapter 4.4 Quantum Numbers Magnetic Quantum Number (m) Orientation of an orbital about the nucleus. l = s m = 0 l = p m = -1, 0, 1 l = d m = -2, -1, 0, 1, 2 l = f m = -3, -2, -1, 0, 1, 2, 3

Chapter 4.4 Quantum Numbers s orbital, 1 orientation.

Chapter 4.4 Quantum Numbers p orbital, 3 orientations. px orbital py orbital pz orbital pxyz orbital

Chapter 4.4 Quantum Numbers d orbital, 5 orientations.

Chapter 4.4 Quantum Numbers f orbital, 7 orientations.

Chapter 4.4 Quantum Numbers Spin Quantum Number(+1/2 , -1/2) Indicates two possible states on an electron in an orbital.

Chapter 4.4 Quantum Numbers Magnetism Caused by the motion of electrons about the nuclei of atoms. Diamagnetism – substance is weakly repelled by a magnetic force. Paramagnetism – substance is weakly attracted by a magnetic force. Ferromagnetism – Strong attraction by a magnetic force.

Chapter 4.4 Quantum Numbers Principal Energy Level Sublevels Orbitals N=1 1s N=2 2s, 2p 2s(one) + 2p(three) N=3 3s, 3p, 3d 3s(one) + 3p(three) + 3d (five) N=4 4s, 4p, 4d, 4f 4s(one) + 4p(three) + 4d (five) + 4f(seven)

Chapter 4.4 Quantum Numbers Principal Q.N. # Orbitals per Main level (n2) # Electrons per main level (2n2) 1 2 4 8 3 9 18 16 32

Chapter 4.4 Quantum Numbers Orbital Max # electrons s 2 p 6 d 10 f 14

4.5 Rules Governing Electron Configurations Electron Configuration – arrangement of electrons in the atom Rules Aufbau Rule – electron occupies the lowest energy level that will receive it. Hund’s Rule – orbitals of equal energy each receive one electron (equal spin) before any receive two. Pauli’s Exclusion Principle – no two electrons can have the same set of 4 quantum numbers (n,l,m,s)

Orbital Notation Orbital Notation Orbital represented by a line ____ Electron is represented by an ½ Arrow  + ½ () - ½ () 

Order of Energy Levels 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d Number - principal quantum number, the main energy level Letter – orbital quantum number, the shape Useful Diagram 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d 7d 4f 5f 6f 7f

Orbital Notation Write the orbital notation for the following elements: Al Zn P Cl 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d 7d 4f 5f 6f 7f

Short-Hand Notation Eliminates the lines & arrows Superscripts are used to illustrate the number of electrons in the sublevel Same order of sublevels

Electron-Configuration Notation Write the electron-configuration for the following: Cs Kr Br Po 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d 7d 4f 5f 6f 7f

Exceptions to Aufbau All elements prefer a more stable configuration of electrons. Fully filled and ½ filled orbitals are more stable than others. Elements that are 1 shy of a full or ½ filled d orbital configuration will have electrons transfer from the s to the d to reach this stable state. Example if you have a 4s2 and 3d4 the actual configuration should be 4s1 and 3d5.

Identifying Electrons Paired electrons – when 2 electrons are within the same orbital. Unpaired electrons – when a single electron is within an orbital.

Identifying Electrons How many unpaired electrons does the following elements have? Na O B