Honors Chemistry Chapter 5

1. Honors Chemistry Chapter 5 Electrons

2. “The more success the quantum theory has, the sillier it looks.” ~Albert Einstein, Nobel Prize in Physics, 1921. “If quantum mechanics has not yet profoundly shocked you, you have not yet understood it.” ~Niels Bohr, Nobel Prize in Physics, 1922. “Quantum mechanics: the dreams that stuff is made of.” ~unknown

3. Where are electrons located? • Outside the nucleus • How are they arranged?

5. Bohr model of electron placement • Called the “planetary” model • Electrons closer to the nucleus – lower “energy level” • Electrons farther away from the nucleus – higher “energy level”

6. Called the “shells” K, L, M, N • 2, 8, 18, 32 electrons • Studied the emission spectrum of Hydrogen • Specific colors that are emitted (given off) when an atom releases energy

7. Quantum Staircase Niels Bohr • An electron in a stable orbit will have a specific, restricted (quantitized) energy:

8. Max Planck • Stated that the object (metal) emits energy in small, specific amounts called Quanta. • Quantum is the minimum quantity of energy that can be lost or gained by an atom. • Step ladder analogy

9. Albert Einstein • Took Planck’s idea a little further. • He introduced that electromagnetic radiation has a dual wave-particle nature. • Light exhibits many wavelike properties • Can also be thought of as a stream of particles

10. Bohr’s Model • Two important concepts from Bohr: • Electrons exist only in certain discrete energy levels • Energy is involved in moving an electron from one level to another • IN REALITY, ELECTRONS DO NOT ORBIT THE NUCLEUS LIKE PLANETS ORBITING A STAR!!!!

11. Electrons not really in “planetary” orbits • Are really in areas of “probability” called “electron clouds” • Quantum Model of electron placement

12. Quantum Mechanical Model • This model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus

13. Heisenburg’s Uncertainty Principle • it is not possible to know both the velocity and position of a particle at the same time • velocity = speed and direction

14. Orbital - 3 dimensional region around the nucleus where a particular electron can be located • “clouds” - that show a region of high probability of finding an electron • size and shape of “cloud” depends on energies of electrons that occupy them

15. Principal Energy Levels • Indicates main energy level of an electron in an atom • called “shells” • 1 = lowest 7 = highest • can be any positive integer

16. Sublevel • Indicates the shape of an orbital labeled s, p, d, f

17. s = sphere • p = dumbbell or figure-eight • d = 4 lobes • f = complicated

21. Orbital shapes for Sc • funky orbitals

22. Principal Energy Levels • Are divided into sublevels • the number of sublevels allowed is equal to the principal energy level (n) (up to n=4) • PEL = 1 1 sublevel • PEL = 2 2 sublevels • PEL = 3 3 sublevels • PEL = 4 4 sublevels

23. Wish you were here?Well, you’re not, so pay attention!

24. “s” sublevel – 1 orbital allowed • “p” sublevel – 3 orbitals allowed • “d” sublevel – 5 orbitals allowed • “f” sublevel – 7 orbitals allowed

25. PEL sublevels allowed 1 s 2 s, p 3 s, p, d 4 s, p, d, f

26. Each sublevel has a certain number of orbitals allowed • Sublevel orbitals allowed s 1 p 3 d 5 f 7

27. Maximum of 2 e- in any orbital ! • They “spin” in opposite directions

28. WAKE UP!!!!!!!!

29. Don’t give up!

30. You can’t escape!

31. I know you’d rather be here, but it gets better, I promise!

32. Are you ready????????

33. Chart that follows this slide: • Principal energy level • type of sublevel • #orbitals per type • #orbitals per level • Max. # electrons

35. Rules for writing electron configurations • Add one electron at a time according to these rules: • 1. each added electron is placed in a sublevel of lowest energy available (Aufbau Process) • 2. No more than 2 electrons can be placed in any orbital (Pauli Exclusion Principle)

36. 3. Before a second electron can be placed in any orbital, all the orbitals of the sublevel must contain at least one electron (Hund’s Rule) • 4. No more than 4 orbitals are occupied in the outermost principal energy level of any atom. (next electron must enter the next principal energy level)

37. NOW!!!! • We will start writing electron configurations • “Regular” and • “Exceptions”

38. Stanford explanation

39. Atoms absorb a SPECIFIC amount of energy – quanta • Electrons “jump” up into energy levels where they really don’t belong • Immediately drop back and release that specific amount of energy in the form of light of specific wavelength and frequency (color)

40. Spectroscopy • Used to study structure of atoms • substances heated • e- move to higher energy levels • “fall back” - release photons of energy of specific wavelength • produce a series of “spectral lines” • are characteristic to specific substances • used as an identifying tool

41. Ground state - atom where the electrons are in the lowest available energy levels • excited state - atom has electrons that have “jumped” to higher energy levels

42. Can identify elements by the colors they produce • FIREWORKS!!! • Flame tests – lab we will do