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Chapter 5

Chapter 5. Electron Configurations and Periodic Trends. True Nature of Waves and Particles. Bohr’s model has some flaws. Bohr’s model has set tracks for electrons to travel, but in reality they were in all space almost at the same time. Like a ceiling fan rotating in three dimensions.

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Chapter 5

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  1. Chapter 5 Electron Configurations and Periodic Trends

  2. True Nature of Waves and Particles Bohr’s model has some flaws. Bohr’s model has set tracks for electrons to travel, but in reality they were in all space almost at the same time. Like a ceiling fan rotating in three dimensions.

  3. Wave-Particle Duality of Nature • Wave-Particle Duality: Light and some forms of matter (like electrons) can be viewed as both a wave as well as a particle. • Created by de Broglie

  4. Heisenburg uncertainty principle Heisenberg uncertainty principle – it is impossible to know both the exact position and the exact momentum of an object at the same time. - the more that we know of one, the less we know of the other. - treats electrons as particles

  5. Schrodinger (ha ha, get your laughs out now….) - tried to treat electrons as waves. - came up with an elaborate equation to describe the location of an electron based upon certain quantum numbers

  6. Waves • As a wave, typical wave properties are observed like wavelength and frequency. • Wavelength – distance from crest to crest - Measured in meters • Frequency – number of cycles in one second • - Measured in Hertz (Hz)

  7. Electromagnetic Spectrum • The electromagnetic spectrum shows a range of all types of electromagnetic radiation (x-rays, gamma rays, visible light, infrared, radio waves) in order of wavelength and frequency. • All types of electromagnetic radiation travels at the speed of light (c = 3.0 x 108 m/s)

  8. Electromagnetic Spectrum

  9. Shortest Longest Wavelength Highest Frequency Lowest Frequency Higher Energy Lower Energy MICROWAVES Communication Electromagnetic Spectrum 5

  10. Waves • Wave speed = Wavelength x Frequency (m/s) (m) (Hz) Ex. 1 What is the frequency of yellow light which has a wavelength of 580 nm? Ex. 2 What is the wavelength of the signal from 1100AM which has a frequency of 1100 Megahertz?

  11. More light examples • What is the frequency of ultraviolet light which has a wavelength of 220nm? • What is the wavelength of gamma rays which have a frequency of 1024 Hz?

  12. Planck’s constant and equation • Planck also noted that there was a relationship between the energy of light and the frequency. The higher the frequency, the greater the energy. • E = h x f • H = Planck’s constant (6.63 x 10 23 Joule seconds) • F = frequency (Hz) • E = Energy (Joules

  13. More examples • How much more energy is present in a gamma ray with a frequency of 1024 Hz compared to infrared rays with a frequency of 1012 Hz? • What is the frequency of electromagnetic radiation that has 65,000J of energy? • How much energy does a ray of green light have if its wavelength is 580nm?

  14. Electromagnetic Spectrum

  15. Emission of Light

  16. When light is given off, the color that you see is based on the elements that are being given energy (from electricity or heat usually). • The electrons are given energy and are “excited” so they move up to the next energy level. • When the electrons return to their “ground state” they give off the energy in the form of light.

  17. We can actually see electrons moving down in energy level through flame tests…energy given off corresponds to the color of the flame produced…

  18. Atomic Spectrum • Each element has its own characteristic spectrum which can be used to identify specific elements by using a spectroscope.

  19. Electromagnetic Spectrum

  20. Quantum Numbers • A quantum is a discrete particle (as opposed to a continuous wave). • Bohr explained the atomic spectrum by saying that the electrons were given energy by a photon of light (not a wave, but a “packet” of light energy) • This photon gave the electron energy as it became “excited”. • The frequency of the color given off was dependent on how much energy the electron absorbed and then emitted.

  21. Quantum numbers Schrodinger proposed four different quantum numbers to describe the location of an electron around a nucleus.

  22. Principle quantum number (n) – corresponds to energy level - goes in order of 1,2,3,…with one being the lowest in energy - n is also used to tell us the distance that an electron is from the nucleus. - based on the period (or row) in periodic table Greatest number of electrons per energy level is given by the equation 2n2. (ex: How many e- in the 1st energy level? - the 2nd? - the 3rd?

  23. Principle Quantum Number • As the principle quantum number increases, the energy level increases and the size of the atom increases.

  24. What are the maximum number of electrons in the second energy level? • 2 • 4 • 6 • 8 • 18 • 32

  25. What are the maximum number of electrons in the fourth energy level? • 2 • 4 • 6 • 8 • 18 • 32

  26. Sublevels - Sublevels - smaller energy states grouped inside a larger energy level • Sublevels in order of increasing energy within an energy level: s < p < d < f Mnemonic tool: “some people don’t fart” or make up your own

  27. Number of types of principal quantum orbitals in energy level = number Energy LevelSublevels 1st energy level 1s 2nd energy level 2s 2p 3rd energy level 3s 3p 3d 4th energy level 4s 4p 4d 4f

  28. How many types of sublevels are in the second energy level? • 1 • 2 • 3 • 4 • 5

  29. How many sublevels are in the fourth energy level? • 1 • 2 • 3 • 4 • 5

  30. Different Shapes of Sublevel Orbitals An orbital is the region that you are most likely going to find a specific electron. Shapes of orbitals: s – spherical p – barbell shaped (each centered on a different axis – x,y,z) -

  31. Orbital – region occupied by one pair(2) of electrons. - s has one orbital = 1pair (2 electrons) - p has three orbitals = 3 pairs (6) - d has five orbitals = 5 pairs (10) - f has seven orbitals = 7 pairs (14)

  32. How many sublevels are in the third energy level? • 1 • 2 • 3 • 4 • 5

  33. Which of the following sublevels is not present in energy level 3? • f • s • p • d • All are present

  34. How many orbitals are in the 2p sublevel? • 1 • 2 • 3 • 4

  35. What is the maximum number of electrons that can fit into the 2p sublevel? • 1 • 2 • 4 • 6 • 8

  36. How many orbitals are in the 4d sublevel? • 1 • 2 • 4 • 5 • 8 • 10 • 14

  37. What is the maximum number of electrons that can fit in the 5p sublevel? • 2 • 4 • 6 • 8 • 10 • 14 • 32

  38. Describe the shape of the 4s orbital? • Spherical • Dumbbell shaped • Cube-shaped • Like a dragon

  39. Overlapping occurs in the third and fourth energy level, as it takes more energy to put an electron into a “d” than an “s” or a “p”… - “d” is filled one energy level later. - “f” is filled two levels later. These shapes are filled later because they are so complex and require additional energy to enter.

  40. Pauli Exclusion Principle - electrons will singly fill each orbital until all orbitals have one atom before putting two electrons in one orbital. - example: filling rooms in a hotel, seats on an airplane

  41. Electron Spin – describes difference between two electrons occupying the same orbital. • Electrons spin in opposite directions when occupying the same orbital • Notation for electrons with differing spins:

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