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Electrons and Quantum Theory

Electrons and Quantum Theory. Electrons around the atom. How do we know how many Electrons are in a particular atom?. In a neutral atom the number of electrons is equal to the: Atomic number # of protons. How many electrons are in a neutral atom of Potassium, K?. 19.

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Electrons and Quantum Theory

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  1. Electrons and Quantum Theory • Electrons around the atom

  2. How do we know how many Electrons are in a particular atom? • In a neutral atom the number of electrons is equal to the: • Atomic number • # of protons

  3. How many electrons are in a neutral atom of Potassium, K? 19

  4. Why learn about the electrons? • Because almost EVERYTHING about how an atom behaves, reacts, and combines is due to its electron arrangement.

  5. Electrons • Can be found in orbitals around the nucleus. • Define an atom’s or compound’s charge • Ion – an atom or group of bonded atoms with a positive or negative charge • 1) fewer electrons than protons = positive (Na+) • 2) more electrons than protons = negative (F-) • question

  6. What’s light have to do with it?A quick Demo!!!

  7. It all has to do with light! • Matter absorbs and emits light because of the arrangement of the electrons.

  8. Remember the Electromagnetic Spectrum? • It is a graph of all electromagnetic waves. • These include TV, radio, microwaves, gamma rays, and more. • These waves all: • Travel at 3 x 108 meters/second

  9. The Electromagnetic Spectrum

  10. Visible light, x-rays, infrared radiation and radio waves all have the same…. • Energy • Wavelength • Speed • Frequency

  11. The higher the frequency, the greater the energy of the wave, the shorter the wavelength, λ.

  12. What is light? • Light is a wave, part of EMS • Light consists of “quanta” or packets of energy • Light consists of a stream of particles known as PHOTONS • Thus, light is both particle and wave

  13. Photon • A particle of the EMS • Has ZERO mass • Carries a quantum of energy • Energy depends on the frequency of the wave

  14. For an electron to change from a ground state to an excited state, • Energy must be released • Energy must be absorbed • Radiation must be emitted • The electron must make a transition from a higher to a lower energy state

  15. White light is made out of colors

  16. Sunlight refracted into its colors

  17. Every atom emits a unique spectrum of light. Hydrogen:

  18. The specific wavelengths of light seen through a prism that are made when a high voltage current passes through a tube of hydrogen gas… • line-emission spectrum • Electron configuration • Photoelectric effect • Continuous electromagnetic spectrum

  19. So why the unique spectrum? • Electrons orbit in different energy levels. • The e- get excited by heat or energy • They then jump to a higher energy level. • When they jump back down, they have to release this energy, often as visible light.

  20. Electron absorbing energy.Electron releasing energy.

  21. A video on the Photoelectric Effect

  22. We can’t see the spectral lines without a special viewing device.

  23. Now you know what causes the color of fireworks; it’s all due to electron jumps!

  24. Do electrons orbit like planets?

  25. The Four Quantum Numbers: • These describe the probable location of each electron in an atom. • Each e- has a unique set of quantum numbers, like a “probable” address. • Instead of number, street, city, and zip, quantum numbers go backwards: zip first, then city, then street, then number.

  26. 1st Quantum Number: nn is the Principal energy level • n=1,2,3,4,5,6 or 7 (corresponds with period of the element) • n determines all other quantum numbers and describes the size and energy of the orbital • It’s like what floor you’re on in a department store.

  27. 2nd Quantum#, the angular momentum Quantum #Sub Level: LLdescribes the shape of the orbital • L = n-1 • L= 0,1,2, or 3 We also use letters: • L =s,p,d, or f • It’s like what department you’re in of the floor you’re on. • n=3rd floor; L=men’s shoes

  28. The three “p” orbitals next to an actual plot of e- locations

  29. 3rd Q# Orbital: mLmL is the orientation of the e- pair • mL = -3,-2,-1,0,1,2, or 3 • Each number stands for just ONE PAIR of electrons. Here’s a “p” sublevel: (↑ ↓) (↑ ↓) (↑ ↓) mL=-1 0 +1 • It’s like the item you’re on. • 3rd floor, men’s shoes, mL = a pair of shoes.

  30. 4th Q# Electron Spin: ms • The electron spins one way or another • up ms= +1/2 • down ms= -1/2 • 3rd floor, men’s shoes, pair of shoes, ms = right shoe

  31. How can the two electrons stay in the same orbital?

  32. Electrons & Quantum Theory • Electron Configurations – A simpler way than drawing and labeling shells/orbitals • A. The Rules • 1) Aufbau Principal - electrons fill lowest energy levels / orbitals first • 2) Pauli Exclusion Principal – electrons in the same orbital have opposite spin (+ 1/2 , -1/2) • 3) Hund's Rule – electrons fill one per orbital before pairing up

  33. Two ways to write electron configurations: • Hunds Orbital Diagrams (arrow diagrams) ____ 1s • Electron Configuration notation 1s2

  34. An easy way to “see” electron configurations: the “blocks” in the periodic table: (a song)

  35. Summary of Quantum the first three quantum numbers:

  36. Write the ground state electron configuration and orbital notation for: • Phosphorous • Nitrogen • Potassium

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