2. The Dilemma  Particles have mass and a specific position in space (matter)  Waves have NO mass and NO specific position in space (light and energy) Is the electron a wave or particle?

3. Wave-Particle Duality Several disputes over whether or not light and the electron were waves or particles. JJ Thomson won the Nobel prize for describing the electron as a particle. George Thomson (JJ Thomson’s son) won the Nobel prize for describing the electron as a wave. The electron is a particle!

4. The Wave Like Electron Louis deBroglie The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves.

5. Types of Waves  Light Waves- vibrations of photons  Sound Waves-vibrations of air molecules and atoms  Water waves-vibrations of water molecules

6. Parts of a wave  Crest: Top of the Wave  Trough: Bottom of the Wave  Wavelength: the shortest distance between equivalent points on a continuous wave (crest to crest/trough to trough) typically measured in meters  Frequency: the number of waves that pass a given point per second. Hertz (Hz) is the SI unit.  Amplitude: the wave’s height from the origin to a crest, or from the origin to a trough. (Not affected by wavelength and frequency)

7. Electromagnetic Spectrum A spectrum that includes all forms of electromagnetic radiation, with the only difference in the types of Radiation being their frequencies and wavelengths. ROYGBIV

8. The wave equation: This speed is constant for all electromagnetic waves inside a vacuum (space) Notice: 1. As frequency increases, wavelength decreases (inverse relationship) 2. As frequency increases, energy of the wave increases Practice: What is the frequency of an X-ray with a wavelength of 1.15 x 10 -10 m?

9. Give it another try: 1. What is the frequency of a water wave that has a wavelength of 5.87m with a speed of 34.2m/s? 2. If a radio wave has a frequency of 8.97 x 10 7 Hz, what is the wavelength of the wave?

10. Warm Up  A sound wave traveling at 350 m/s has a frequency of 500 Hz. What is its wavelength?

11. Is light a particle or wave?  Light as a wave failed to explain: 1. Why heated objects emit only certain frequencies of light The colors produced correspond with different wavelengths and frequencies. 2. Why some metals emit electrons when light at a given temperature shines on them (photoelectric effect)

12. Max Plank (1858-1947)  While studying the radiation emitted by solid bodies heated, known as black body radiation, he discovered that some how light and matter had to mingle.

13. Plank’s Theory (1900)  There is a fundamental restriction on the amount of energy that an object emits or absorbs Matter can either gain or lose energy but only in small specific amounts called quanta Quantum- is the minimum amount of energy that can be gained or lost by an atom. ○ E= h x ʋ E – energy ʋ - frequency h – Plank’s Constant (6.6262 x 10 -34 J/s)

14. Plank’s Theory (1900)  Think of the dots as stepping stones: Each energy level is a stepping stone Electron transitions involve jumps of a definite amount of energy Each transition produces bands of light with definite wavelengths. (Specific color) As an excited electron returns to the ground state they emit energy, that appears as specific colors of the specific energy levels.

15. Photoelectric Effect:  Some metals will eject electrons from their surface when light of a certain frequency (or higher) hits their surface.  Solar Power

16. Albert Einstein  Duality of Light: Light can be both a wave and a particle  Proposed that light consist of quanta of energy that behave like tiny particles (photons) Photons- a massless particle that carries a quantum of energy. This energy depends of the frequency of the photons.

17. Neils Bohr (1913)  Studied the hydrogen atom  determined that the atom only had certain allowable energy states Ground State - lowest possible energy state Excited State - when the atom absorbs energy

18. Neils Bohr  Bohr suggested that the electrons around the hydrogen atom could only be allowed in certain circular orbits around the nucleus The smaller the electron’s orbit, the lower the atom’s energy state or energy level The larger the electron’s orbit the higher the atom's energy level  Quantum Number: the number Bohr gave to each orbital around the atom

19. Energy State  Ground State for hydrogen is 1s 1 (n=1) Meaning H has a single electron in the first energy level H does not give off energy in the ground state  When energy is added, a single electron moves up to a higher energy level creating an excited atom The electron will fall back into its original quantum level (ground state) and release the energy it gained as a photon (color)  Only specific frequencies are emitted by an atom Quantum's only allow a certain amount of energy to be absorbed and emitted by the atom

20. Atomic Emission Spectrum  Electrons absorb energy, jumping up to a different energy level, and release the same amount of energy when falling back to the ground state When the electrons falls back to the ground state is when it releases the photon, a specific color is then seen.  The amount of energy has a specific frequency, that is visible in colors The color is the photons being released, and they are within the visible light of the electromagnetic spectrum  Atoms absorb then release energy in the form of light Every element emits light containing only certain wavelengths Each element has a very specific range of colors that are emitted.

22. Wave-Particle Duality

23. Louise de Broglie (1924)  Thought that if light can have both wave and particle properties/characteristics, then so could matter (electrons)  Predicted that all moving particles have wave-like characteristics

24. Werner Heisenberg (1901-1976)  Stated that it is impossible to take any measurement of an object without disturbing the object  Heisenberg Uncertainty Principal: states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time *Meaning that Bohr’s defined orbits were not accurate

25. Erwin Schrodinger (1926)  Quantum Mechanical model of the atom: the atomic model in which electrons are treated as waves This allowed scientist to determine particular volumes of space around the nucleus in which the probability of finding an electron is very high  Atomic Orbital: the probable location of an electron within the atom

26. Answering the Dilemma of the Atom  Treat the electron as waves  As the electron moves away from the nucleus, the wavelength shortens  Shorter wavelengths = higher energy  Higher energy = greater distance from the nucleus