Electromagnetic Radiation & Light

2 What are the atom models we know of? 2

Models of the Atom A. There were many different models over time 1. Dalton-billiard ball model (1803) 2. Thompson – plum-pudding model (1897) 3. Rutherford – Nuclear model of the atom (1911) 4. Bohr – uses quantized energy of the atom (1913) 5. Quantum Mechanical Model of the Atom (1926)

4 How do the atom models show the importance of communication and sharing information in the sciences? 4

B.Each new model contributed to the model we use today. Even our current Quantum Mechanical model, does not give us an exact model of how electrons interact.

6 Why would you use the simplest for of an atom to make a model? What is the simplest atom? What is another name for a circular pathway? What would be the difference if you were in a grounded state vs an excited state? How could that relate to electron orbitals? 6

The Bohr Model of the Atom A. Bohr used the simplest element, hydrogen, for his model B. He proposed an electron is found in specific circular paths, or orbits around the nucleus C. Each electron orbit was thought to have fixed energy levels D. Lowest level-ground state E. Highest level-excited state

8 What could quantum mean? What would we need to do to excite an electron? 8

F. One electron is capable of many different excited states G. An e- can gain or lose is a specific amount or quantum of energy H. You can excite an e- with energy like electricity, the sun, or magnets

10 What is released when an electron goes to a lower energy level? What are photons? 10

Explanation of Atomic Spectra step1 step2 When an e- is excited it moves to a higher energy level A.When the e- drops back down to a lower energy level, it gives off a quantum of energy called a “photon” Only certain atomic energies are possible and emitted Photons have both wave like and particle like properties E.When a photon is emitted, we emit energy, and use the equation E = h  ν

Energy levels are like rungs of a ladder. You cannot be in between a rung Energy levels in an atom’s electron are unequally spaced. The higher energy levels are closer together.

13 What does spectrum mean? Give an example What would the electromagnetic spectrum mean? 13

Electromagnetic Spectrum A. Electromagnetic spectrum is the range of all energies emitted from photons acting like waves.

Electromagnetic Spectrum with Visible Light Spectrum

16 What do these terms mean when related to a wave? 16 Wavelength Frequency Amplitudetrough Crest

Characteristics of a Wave Amplitude (Wavelength) A. Wavelength  (lambda) – shortest distance between equivalent points on a continuous wave [Unit = meters] B. Frequency  (nu) – the number of waves that pass a given point per second [Unit = 1/second = s -1 = Hertz (Hz)] C. Crest – Highest point of a wave D. Trough – Lowest point of a wave E. Amplitude (a)– height from its origin to its crest (highest point) or trough (lowest point) [Unit = meters]

18 What is the relationship between wavelength and frequency? 18

Wavelength and Frequency A. Wavelength (  ) and frequency (  ) are related B. As wavelength goes up, frequency goes down C. As wavelength goes down, frequency goes up D. This relationship is inversely proportional

c =  Speed of light (c) = 3 x 10 8 m/s c  wavelengthfrequency Speed of light c =   = c /   = c / 

Question Time  c What is the frequency (  ) of radiation with a wavelength (  ) of 5.00 x m? What region of the electromagnetic spectrum is this radiation? How do we calculate the energy of a wave? What are the units for energy? How are energy and frequency related? What happens when an electron goes to a lower energy level? How much energy is given off by a wave with a frequency of 2.0 x 10 8 Hz? ( h = Planck’s constant = x Joule. Sec)

22 Which colors of light have the most/least energy? Is the high energy light going to have a higher or lower frequency? Wavelength? 22

Visible Light, Frequency, and Energy A.Red 1. longest wavelength (  ), 2. smallest frequency (  ) 3. least amount of energy (E) B.Violet 1. smallest wavelength (  ), 2. largest frequency (  ) 3. greatest amount of energy (E)

Flame Test The flame test is a way to determine the certain elements that are present in a sample When placed in a flame, different elements give off different colors Why would they give off different colors? If no colors are emitted could there still be photons released? How?

Why Do We See Different Colors in the Flame Test? A. The electrons are energized to an excited state B. As electrons drop to lower levels, they give off photons C. A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy D. If the photon’s frequency corresponds the visible light range we can see different colors in the flame E. If it is not in the visible light range, it may be giving off other forms of electromagnetic radiation like radio, microwaves, infrared, ultra violet, x-rays, or gamma waves

26 How could we use the flame test/electromagnetic spectrum? 26

Atomic Emission Spectrum A. When a current is passed through a vacuum tube of gas at low pressure, a set of frequencies of the electromagnetic waves are emitted by atoms of the element B. Used to determine which elements are present in a sample C. Used to determine which elements are present in a star D. Each element has a unique spectrum E. Only certain colors are emitted meaning only certain frequencies of light are emitted

Spectroscope A. A spectroscope is needed to see the atomic emission spectra, which acts similar to a prism, separating different frequencies of light

Problems with the Bohr Model A. The model is fundamentally incorrect and only works with hydrogen. B. Failed to explain the spectrum of any other element C. Did not account for the chemical behavior of atoms D. Electrons do not move around the nucleus in circular orbits E. Laid ground for later atomic models.

Emission Spectrum for Hydrogen Absorption Spectrum for Hydrogen Problems with the Bohr Model cont.