2. Split Personality

3.  Electrons are negatively charged particles that move around the positive nucleus.  Why don’t they get “sucked in” by the positive charge?  The energy associated with their motion keeps them outside the nucleus!

4.  energy created when the e - “wiggles” making waves as it moves  particles which can absorb little bundles of energy called photons

5.  DeBroglie called this Wave-Particle Duality since e - can move in a wave-like motion and act as particles capable of absorbing photons! 1892-1987 e - orbit forms a wave around nucleus as it moves

6. A wave is just the way the energy moves. Wavelength (λ and measured in meters)= crest to crest or trough to trough Frequency (ν and measured in Hertz)= number waves that pass a point in a measured period of time wavelength

8. Wavelength Variances  waves close together = more frequent waves = more energy  waves far apart = less frequent waves = less energy

9. Which wave below has MORE ENERGY?

10. If the λ increases, the ν decreases, and if the ν increases, the λ decreases because all types off electromagnetic radiation travel at the speed of light! C=λν

11. RRemember: Rutherford’s Nuclear Atomic Model organized the nucleus. SScientists began to wonder how the e - outside the nucleus are arranged. BBohr came up with the Planetary Atomic Model which was: ee - move around the nucleus in circular orbits (paths) much like planets orbit the sun

12. Bohr’s Planetary Atomic Model

13.  Each “orbit” of an e - has a certain level of energy.  So...we call the orbits/paths of e - energy levels (which are in e - cloud). smaller orbits = lower energy larger orbits = higher energy  The e - now is acting particle like...

14.  The LOWEST energy level in e - cloud is called the ground state.  The ground state is the orbit/path closest to the nucleus.  The ground state is the most stable energy level. It’s the “home-base” of an e -.

15. ground state higher energy level nucleus

16. Electrons are able to change energy levels (like climbing the rungs of a ladder). An e - can change energy levels by absorbing little bundles of energy called photons

17.  “Jumping” to higher energy level = absorb a photon = e - moves lower to higher energy level = excitation Wouldn’t YOU need energy to “jump” up high?! Well...so does an electron!

18.  Once an electron is “excited” and in a higher energy level, it can’t stay there because it’s not the “home” orbit.

19.  “Excited” electron = emits photon (excess energy) = drops higher to lower energy level = de- excitation = see visible light  The e - releases the photon before it “drops” from the higher energy level to the lower energy level. The photon emitted by e- during de- excitation corresponds to light in the visible electromagnetic spectrum. Electron Orbits

20. RRed light is lowest energy; longest wavelengths. VViolet light is highest energy; smallest wavelengths.

21.  Neon lights give off a red-orange color?  Some fireworks are purple? Others white? And some all different colors?

22. Have you ever wondered...  How we know what elements are in the sun and stars?  How we know what elements are in comets? This is all because of the excitation and de- excitation of e - as they absorb and release photons (energy)!

23. Are you still awake??

24.  How did scientists figure all this out?  Scientists noticed when certain elements are burned, they emit visible colored light! Why?  e - absorbed the heat energy (photons) from the flame.  The heat energy causes the e - to excite.  The e - then emits a photon during de- excitation.  The photon emitted corresponds with the same energy as a wavelength of a certain color.

25.  The light given off when an element is heated is called it’s atomic emission spectrum.  This is a discontinuous band of colors (interrupted by black) that represent light in the visible part of the EM. These lines are called SPECTRAL LINES

26.  Every bar of color represents de-excitation of an electron.  How many times did the e - change energy levels (or de-excite)?  20 – b/c there’s 20 different lines each with a wavelength of a different color.

27.  The atomic emission spectrum is like the “signature” of that element – it’s a physical property which is different for every element!  It can be used to identify the element like a fingerprint can identify you!

28. No two elements have the same atomic emission spectrum!!!

29. Heisenberg’s Uncertainty Principle It is imposible to know the exact position and velocity of an electron as it moves around the nucleus. We can know it is in a certain energy level or moving to a higher or lower one, but not where it is EXACTLY or how fast it is moving.

30. There is a GOOD probability of finding an electron in an atomic orbital. An atomic orbital is a 3-D region around the nucleus (it’s not an orbit – contrary to what Bohr thought).

31. 1. Developed by Schrödinger. 2. Atoms have a dense, positively charged nucleus. 3. Electrons surround the nucleus and are treated as waves. 4. Electrons “live” in the 3-D atomic orbitals that lie within energy levels of the e - cloud.

33. Let’s Practice! Just kidding!

34. Was it the worst day of your life?