1. Ch. 5 Electrons in Atoms This Quiz = Begin learning/memorizing the elements and symbols for 1-20,22,24-30, 33,35,36,38 Next Quiz = Learn 47,50,51,53,54,78-80,82,92

2. Ch. 5 Electrons in Atoms Fluorine #9 = most reactive element and a gas Neon #10 = 2nd most stable/non-reactive element and a gas Sodium #11 = one of the most reactive elements and a solid KC #1: Rutherford/Bohr Model was too simplistic and couldn’t explain the huge range of varied properties found in atoms. Extreme changes in properties meant complex structure

3. Ch. 5 Electrons in Atoms Bohr Model expanded on Rutherford by proposing that energy in atoms comes in “increments” and electrons are found on Energy Levels

4. Bohr Model of the Atom In 1914 Niels Bohr proposed that the energy levels for the electrons in an atom are quantized E n = -hcR H (1/n) 2 E n = (-2.18  10 -18 J)(1/n 2 ) Where n = 1, 2, 3, 4, … n=1 n=2 n=3 n=4

5. Ch. 5 Electrons in Atoms Light and energy absorption and emission Same idea, just with particle accelerators and electron microscopes The specific energies that an electron can have The amount of energy absorbed to move an electron from one energy level to another No, in fact they are different for every atom, regardless of energy levels Worked well for explaining the behavior of hydrogen, but not other elements (F, Ne, Na)

6. Ch. 5 Electrons in Atoms A set of mathematical equations that described the range of possible energies an electron could have Determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus A mathematically defined region based on a % likelihood of find an electron with the given area

7. Ch. 5 Electrons in Atoms The energy levels around an atom (1-7) The divided regions of space within an Energy level (spdf) Each energy sublevel corresponds to an orbital (probability cloud) of a different shape (all ‘s’ are alike) s p d f, shape and amounts of energy s=2p=6 d=10f=14 use the formula above and the periodic table

8. Ch. 5 Electrons in Atoms Let’s stop for a minute, and review…. 1. Electrons are found around the nucleus in an extremely complex organization 2. This organization is predicted and described by a set of math equations called Schroedinger’s Equations. 3. The complex organization of the electrons explains and predicts the properties and behavior of elements

11. E. L.Sublevels Orbitals Electrons 1s1 2 2s p1 3 2 6 3spd1 3 5 2 6 10 4s p d f1 3 5 7 2 6 10 14 5s p d f “g” 6 7

15. Noble Gas Shorthand

16. Orbital Diagrams

17. Ch. 5 Electrons in Atoms - The arrangement of electrons of an atom in its ground state (no absorbed energy) into levels and orbitals - Some elements “break the rules” like Cu and Cr on page 136 Give an example of an electron configuration. 1s 2 2s 2 2p 6 3s 2 3p 6 Give an example of a Noble Gas shorthand configuration. [ Ne ] 3s 2 3p 5 Given an example of an orbital diagram. _  _     _ 1s 2s 2p

18. Ch. 5 Electrons in Atoms Ch. 5.2 - Aufbau, Pauli, Hund Ch. 5.3 - Energy waves that travel at the speed of light but have differing wavelength and frequency - The range or spectrum of electromagnetic radiation as they gradually change wavelength and frequency

19. λ = wavelength a = amplitude ƒ = frequency (waves/s) v = speed of light 300,000,000 m/s

21. Electromagnetic Spectrum

22. Light as a Wave - Inversely proportional - They would “tie” as they all travel the speed of light c = c = = wavelength (m) = frequency (s -1 ) c c = speed of light (3.00  10 8 m/s)

23. Ch. 5 Electrons in Atoms - When atoms absorb energy, electrons move into higher energy levels. These electrons lose energy by emitting light when they return to lower energy levels - The lowest possible energy of an electron and the energy level it moves to after absorbing energy - The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electrons

24. Blackbody Radiation & Max Planck The classical laws of physics do not explain the distribution of light emitted from hot objects. photons Max Planck solved the problem mathematically (in 1900) by assuming that the light can only be released in “chunks” of a discrete size (quantized like currency or the notes on a piano). We can think of these “chunks” as particles of light called photons. E = h E = h E = hc/ E = hc/ = wavelength (m) = frequency (s -1 ) h h = Planck’s constant (6.626  10 -34 J-s)

25. Line Spectrum of Hydrogen In 1885 Johann Balmer, a Swiss schoolteacher noticed that the frequencies of the four lines of the H spectrum obeyed the following relationship: = k [(1/2) 2 – (1/n) 2 ] = k [(1/2) 2 – (1/n) 2 ] Where k is a constant and n = 3, 4, 5 or 6. n=3n=4n=5n=6

26. Ch. 5 Electrons in Atoms - The energies emitted by an Hydrogen atom that fall in the visible portion of the spectrum - Atomic emission spectra are the patterns formed when the unique shade of light generated by a certain element passes thru a prism and is separated into different freq it contains - Again, Bohr could explain Hydrogen with a single electron, but not the complexities of more electrons - As a particle. - Almost exclusively as a wave

27. Ch. 5 Electrons in Atoms - That energy came in quantums or “packets” - A “particle” of energy - Waves and particles were seemingly opposites, yet light exhibits both characteristics

28. Photoelectric Effect In 1905 Albert Einstein explained the photoelectric effect using Planck’s idea of quantized photons of light. He later won the Nobel Prize in physics for this work.

29. Ch. 5 Electrons in Atoms - Experiment performed in multiple ways, but classically as red and violet light shined on sodium metal - The red light never “worked” to eject electrons and the violet light always “worked” - Proved that energy has particle characteristics -The red light did not provide the “right” amount of energy to make the electrons move an energy level

30. Louis DeBroglie & the Wave- Particle Duality of Matter While working on his PhD thesis (at the Sorbonne in Paris) Louis DeBroglie proposed that matter could also behave simultaneously as an particle and a wave. = h/mv = h/mv = wavelength (m) v v = velocity (m/s) h h = Planck’s constant (6.626  10 -34 J-s) Only important for matter that has a very small mass (electrons). As mass goes up wavelength goes down. We will see later in some ways electrons behave like waves.

31. Electron Diffraction Transmission Electron Microscope Electron Diffraction Pattern

32. Heisenberg & the Uncertainty Principle While working as a postdoctoral assistant with Niels Bohr, Werner Heisenberg formulated the uncertainty principle.  x  p = h/4   x  x = position uncertainty  p  p = momentum uncertainty (p = mv) h h = Planck’s constant We can never precisely know the location and the momentum (or velocity or energy) of an object. This is only important for very small objects. - Classical Mechanics adequately describes the motions of larger bodies, while quantum mechanics describes the motions of subatomic particles and atoms and waves

33. Ch. 5 Electrons in Atoms - Einstein theorized that if we ever perfectly froze an atom, it would violate Heisenberg, so it would instead form a an entire atom probability cloud - Uncertainty coupled with duality of light and matter led to Schroedinger’s equations

34. Schrodinger and Electron Wave Functions Erwin Schrodinger, an Austrian physicist, proposed that we think of the electrons more as waves than particles. This led to the field called quantum mechanics. In Schrodinger’s wave mechanics the electron is described by a wave function, . The exact wave-function for each electron depends upon four variables, called quantum numbers they are n = principle quantum number l = azimuthal quantum number m l = magnetic quantum number m s = spin quantum number

35. Ch. 5 Electrons in Atoms - These equations led to the idea of “quantum numbers” or what we generically refer to as electron configuration. - Further study of this concept through an experiment called the “Double Slit Experiment” determined that electrons also exhibited wavelike properties. This led to the Idea of Quantum Mechanics and the theory or energy absorption and emission.

36. Ch. 5 Electrons in Atoms 1. At atom’s electrons actually absorb the energy 2.The electrons absorb a photon of just the right amount of energy 3.The just right amount of energy is called a quantum 4.The electron is on an energy level called its ground state 5.It absorbs the photon containing a quantum of energy, then jumps to its excited state 6.The electron is now energized and unstable 7.The electron releases the same amount of energy and goes back to its ground state 8.The energy is released typically in the form of visible light or color 9.Repeats

37. s-orbital Electron Density (where does the electron spend it’s time)  2 = Probability density # of radial nodes = n – l – 1

38. Velocity is proportional to length of streak, position is uncertain. Position is fairly certain, but velocity is uncertain. Schrodinger’s quantum mechanical picture of the atom 1. The energy levels of the electrons are well known 2. We have some idea of where the electron might be at a given moment 3. We have no information at all about the path or trajectory of the electrons

39. s & p orbitals

40. d orbitals # of nodal planes = l

41. Electrons produce a magnetic field. All electrons produce a magnetic field of the same magnitude Its polarity can either be + or -, like the two ends of a bar magnet Thus the spin of an electron can only take quantized values (m s =+½,-½), giving rise to the 4th quantum number

42. Single Electron Atom Multi Electron Atom