Chemistry I Honors--Unit 3: Quantum Mechanical Model of the Atom & Periodic Trends

1. Chemistry I Honors--Unit 3: Quantum Mechanical Model of the Atom & Periodic Trends

2. Objectives #1-7: The Development of a New Atomic Model • I. Electromagnetic Radiation • In the late 1800’s and early 1900’s, scientists discovered that passing an electric current through gases of various elements caused electromagnetic radiationin the form of colored lightto be emittedfrom the gas.

3. Examples of colors produced by the electrically charged gasses include:

4. Additional testing showed that EMR of energies too low or too high to see with the eye were also produced • Electromagnetic radiation is energy that travels in the form of a wave

6. Examples of Electromagnetic Waves All waves have AMPLITUDE , FREQUENCY & WAVELENGTH

7. Characteristics of EMR: • The wavelength of a wave is the distance between the peaksof the wave • The frequency of a wave is the number of peaks that pass by a point is space in one second (the rate of reproducibility) • The speed of all EMR is the speed of light(3.00 X 108 m/s)

8. Wavelength and frequency are inverselyrelated to each other c = (f) () • Frequency and energy are directly related to each other E = (h) () • Wavelength and energy are inversely related to each other E = (h)(c) / () h= 6.626 x 10-34 J . Sec These relationships were discovered by Max Planck.

9. Types of EMR Lower Energy Higher Energy Radio Radar Micro IR Visible* UV X-rays Gamma red, orange, yellow, green, blue, indigo, violet

10. Objectives #1-7: The Development of a New Atomic Model • Electromagnetic Radiation * EMR refers to all the various types of radiant energy, from radio waves to gamma waves

11. II. The Origins of Wave MechanicsEMR has dualqualities: (Louis DeBroglie, 1892-1987), French “EMR is like 4 year old who can’t make up their mind what to be for Halloween!! “A wave! No, wait, a particle! No wait, a wave… no wait…..”

12. EMR acts as a particle when it interacts with matter; this is illustrated by the photoelectric effect which involves the emission of electrons when radiation of a specific frequency strikes the surface of a metal(Albert Einstein, 1905, German-American)…  = (h) (v0) • EMR acts as a wave when it travels through space… •  = (h) /(m) (v)

13. Photoelectric Effect Incoming energy waves Electrons emitted Video Clip: The Photoelectric Effect

14. Albert Einstein, 1905, German-American “The photoelectric effect helps explain why your solar calculator works!”

15. Wave Particle Duality • De Broglie discovered that EMR acts as a wave when it travels through space… •  = (h) /(m) (v)

16. The Double Slit Experiment • Evidence of wave particle duality of electrons… • Video Clip: “The Infamous Double Slit Experiment”

17. Objectives #1-7 The Development of a New Atomic Model • Atomic Emission Spectrum: an explanation of the colors produced by exciting atoms • Niels Bohr, 1885-1962, Danish “Electrons LOVE to jump energy levels —they emit light as they move back ‘home’!!”

18. Examples of Atomic Emission Spectrums

19. Bohr’s Theory of Light Emission • An electron is normally in its low energy state or ground state. • When the electron becomes excitedwith a certain amount of energy or quantum, it will “jump” to a higher level of energy or its excited state. • This new state is unstable for the electron and so this excess energy is emitted as a photonof EMR and the electron returns to the ground state.

20. The Bohr Model of the Atom • Proposed that electrons revolve around the nucleus in definite paths or orbits • Each electron has a certain amount of energy associated with it • Electrons are confined to specific energy levels • In order to move from one level to the next, an electron must absorb or release a certain quantumof energy

21. The Bohr Model and Electron Transitions Illustrated A QUANTUM of energy is absorbed to jump levels from GROUND STATE to EXCITED STATE Returning to GROUND STATE , the electron releases energy as PHOTON(S) of light! The color of the light you see is related to the amount of energy being released!!

22. Operation of Spectroscope

23. The Bohr Model Let’s draw some Bohr diagrams!!!!

24. 11Na • 1H • 2He

25. IV. The Quantum Mechanical Model of the AtomErwin Schrodinger, 1887-1961, Austrian “Hmmm…Matter also has particle and wavecharacteristics. So since matter is made of atoms, and atoms contain electrons, then electrons could also travel in waves!!THAT’S IT!!! I’ll be famous!! Won’t Mom be proud!!”

26. Proposed the quantum mechanical model for electrons. • The exactpath of the electron can not be determined because it is traveling near the speed of light and is too small in size. • This idea was based on the work of Werner Heisenberg, 1901-1976, German. • In the Heisenberg Uncertainty Principle there is a limit to how certain we can be about the position and speed of very tiny particles such as electrons.

27. Werner Heisenberg, 1901-1976, German “Its impossible to know the position and the speed of an electron at any given moment—kind of like trying to see Road Runner’s legs when he’s running from Wylie Coyote!”

28. Heisenberg Uncertainty Principle Where’d it go?? Where’d it go???

29. In the quantum model only the probability of finding the electron in a certain area can be determined. • The most highly probable location for an electron about the nucleus is the orbital. • The combination of these areas about the nucleus is called the electron cloud.

30. Objectives #8-10: Quantum Numbers • Schrodinger’s Equation: EΨ = -h2/2m(ð2Ψ/ðx2 + ðΨ/ðy2 + ð2Ψ/ðz2) + V(x, y, z)Ψ (Neat, huh? No you do NOT have to use it, memorize it or solve it!!! You’re welcome! But it IS cool… ) • Solving the previous equation produces various orbital shapes, just as solving y = 1/2x + 2 produces a straight line.

31. Quantum Numbers • Describe energy and location of electrons • Every electron in an atom is unique; each electron has a different energy and therefore will have a different set of quantum numbers.

32. Evolution of the Bohr Model into the Quantum Mechanical Model The energy level number is equal to the number of subshells within that energy level…

33. Principle Quantum Number (n) • Indicates energyand distance from nucleus • Indicates energy level number • Can take on values of: 1 infinity, but 1 7 is currently verified & understood. • Orbital (Angular Momentum) Quantum Number (l) • Indicates shape of orbital (sublevel) • Can take on values of: 0  n-1 (0,1,2,3,etc) • Orbitals that have the same value of n, butdifferent l values are in sublevels,which are designated by letters to avoid confusion!

34. Possible Orbital Shapes:

35. Orbital Shapes s = 1 direction p = 3 directions d = 5 directions f = 7 directions

36. The number of orbital (sublevel) shapes in a level is equal to the level number:

37. C. Magnetic Quantum Number (ml) • Indicates the orientation (direction) of the orbital in space • Indicates the number of orbital directions in a sublevel • Can take on values of: –l  0  +l

38. C. Magnetic Quantum Numbers (ml)

39. Spin Quantum Number (ms) • Indicates the direction of electron spin • Can take on values of: +1/2, -1/2 • No more than 2 electrons can occupy a single orbital • Problems Involving Quantum Numbers • See notes

40. IV. Summary of Electron Energy Level Capacities (See Chart in Lecture Guide) • Some relationships to notice: • If “n” is the number of levels, then the number of sublevels is equal to “n” • If “n” is the number of levels, then the total number of orbitals in a level is equal to n2 • If “n” is the number of levels (and every orbital can hold up to 2 electrons), then the total number of electrons in a level is equal to 2n2

41. Objectives #11-12: Electron Configurations • Electron configurations show electron arrangement • Rules Governing Electron Configurations • The Aufbau Principle • Electrons enter orbitalsof lowest energy first • The Pauli Exclusionary Principle • An atomic orbital may describe, at most, 2electrons

42. Hund’s Rule • Electrons enter orbitals of the same energy with the samespin until each orbital contains one electron before pairing begins • Examples of Electron Configurations w/Orbital Notations & Noble-gas Configurations • 1H: • 2He: • 3Li: • 6C: • 15P:

44. The Diagonal Rule Used to track the order of electrons as they fill available orbitals, according to the Aufbau Principle. Remember that all electrons fill using the LOWEST amount of energy possible!

45. More Examples…Use the Diagonal Rule!19K:___________________________26Fe:__________________________31Ga:__________________________38Sr:_________________________54Xe:_________________________

46. Noble Gas Configurations: Write the symbol of the closest noble gas that is LOWER in atomic number. Then, write out the remaining part of the configuration, following the diagonal rule. Example: 74W—Tungsten Exceptions to the Aufbau Principle: 24Cr: 29Cu:

47. Objectives #13-20: The Periodic Table & Periodicity of Properties • Development of the Periodic Table • The work of Mendeleev (1871, Russian) • Elements were grouped by their properties; • Allowed for prediction of new elements; • Elements arranged by increasing atomic mass

48. Dimitri Mendeleev,1834-1907, Russian “I ALMOST have it…if I could just find those pesky missing pieces!!”

49. Dimitri Mendeleev’s Periodic Table - 1871

50. Henry Moseley,1887-1915,English “Hmmm… what if we used atomic NUMBERS rather than mass…”