1. Chapter 7 Atomic Structure

2. Electromagnetic Radiation Light is a form of electromagnetic (EM) radiation –All forms of EM radiation are types of kinetic energy –7.1

3. EM Radiation Describe each form of EM radiation by its: –Wavelength –Frequency –Energy of its photon

4. All forms of EM radiation travel at the speed of light (c) Wavelength x frequency = speed of light

5. EM Radiation The longer the wavelength the: –lower the frequency and the lower the energy of the EM radiation The shorter the wavelength the: –higher the frequency and the higher the energy of the EM radiation

6. EM Radiation Compare the wavelength, frequency, and energy of: –Ultraviolet light and infrared light

7. EM Radiation The higher the energy of the EM radiation the more damaging it is to living tissue. –CONSIDER GAMMA RAYS, X-RAYS, AND UV LIGHT

8. Planck’s Constant Use Planck’s constant and to calculate the energy of a photon of EM –The smallest “packet” or “quantum” of energy associated with a given wavelength of EM radiation. 7.2

9. Emission of Energy by Atoms Thanks to the work of Planck and Bohr we know that: –When atoms are energized by an input of energy their electrons are excited (energized) –When excited electrons return to lower energy states they emits energy in the form of light. Emits photons of energy Energy of the photons emitted depends upon how excited the electron was.

10. Photoelectric Effect Electrons will leave a metal when light of sufficient frequency strikes the metal. –Called the threshold frequency Light with frequency < threshold frequency – no electrons emitted

11. Photoelectric Effect Light with frequency > threshold frequency –Number of electrons given off increases with intensity of the light –Kinetic energy of the electrons emitted increases as the frequency of the light increases Consistent with EM radiation being quantized After much math this lead Einstein to…. E = mc 2

12. Planck and Einstein Energy is quantized –Occurs in discrete packets called quanta Similar to $ comes in discrete quantities, penny, nickel, dime, quarter EM radiation has both wave properties and particle/matter properties –Each wavelength is associated with a specific quantum of energy

13. Bohr Experiment (~1911) Bohr excited hydrogen atoms by running electricity through a tube of hydrogen gas. –The gas gave off a pink light.

14. Bohr Experiment (~1911) Bohr aimed a beam of the pink light at a “prism” –Found the pink light generated a line spectrum not a continuous spectrum Line spectrum – specific colors of light observed Continuous spectrum – all colors of light present

15. Bohr Experiment (~1911) He observed 4 bands of color: –Purple (410 nm) –Blue (434 nm) –Green (486 nm) –Red (656 nm) –7.3 and 7.4

16. Bohr Experiment (~1911) He then calculated the energy of each color of light –_____________ was the highest energy and _____________ was the lowest energy.

17. Bohr’s Interpretation of the Data Bohr proposed: –The electrons circle the nucleus in orbits of specific energies. –Electrons are always in one of the circular orbits. –Larger orbits are of higher energy than smaller orbits

18. –The electricity excites electrons and allows them to move to higher energy orbits. –When the excited electrons return to lower energy orbits they emit energy in the form of light. –Because 4 specific wavelengths of light are emitted by hydrogen there must be 4 possible orbit changes –For each orbit change, the difference in energy between the orbits corresponds to the energy of the light emitted.

19. Bohr Model of the Atom

20. Calculating Energy Changes for Hydrogen See equation in section 7.4 and example 7.5

21. Bohr’s Model When Bohr’s mathematical approach was applied to other elements it didn’t work. –Bohr’s model of the atom has been revised to replace the circular orbits with “wave mechanical model” of the atom

22. Modern Atomic Structure Still picture electrons to be at specific energy levels, but no longer picture them as traveling in circular orbits. –The current model of the atom locates electrons in orbitals.

23. Orbitals Each orbital is of a specific energy, size, and shape Each orbital can hold a maximum of 2 electrons of opposite spin (Pauli exclusion principle)

24. Orbitals The exact path of an electron in an orbital is not known. –Heisenberg uncertainty principle states that it is impossible to determine the location and path of an electron at the same time

25. Orbitals Orbital shapes describe the region in space where an electron will be found 90% of the time. Each orbital is described by 3 quantum numbers….and each electron by 4 quantum numbers

26. Modern Atomic Theory Atoms have specific energy levels in which electrons may be found. –Called Principal Energy Levels (PEL) –PEL farther from the nucleus are larger and of higher energy. –Assign a number (n) to each PEL See board

27. Modern Atomic Theory Within each PEL are sublevels –Sublevels are named: s, p, d, and f –The larger the PEL the more sublevels it contains

28. Modern Atomic Theory PEL, n# sublevelsTypes of sublevels 11S 2 3 4

29. Modern Atomic Theory PEL, n# sublevelsTypes of sublevels 11s 22s, p 33s, p, d 44s, p, d, f

30. Modern Atomic Theory Sublevels contain orbitals. Sublevel# orbitalsMax. # electrons s1 p3 d5 f7

31. Describing Orbitals See pages 7.7 for diagrams of the orbitals –S orbitals are spherical –The 3 p orbitals are shaped

32. PELSublevels# e per sublevel Max. # e per PEL 1s 2spsp 3spdspd 4spdfspdf

33. PELSublevels# e per sublevel Max. # e per PEL 1s22 2spsp 2626 8 3spdspd 2 6 10 18 4spdfspdf 2 6 10 14 32

34. Putting it All Together Unless they are excited, electrons always occupy the lowest energy orbital with room. –Electrons enter orbitals of a given sublevel one at a time before pairing up (Hund’s rule) –Consider 2 electrons in a p sublevel:

35. The fun part! Our goal is to write the following for atoms and ions: –Electron configuration –Box/energy diagram –Lewis dot symbol Our goal is also to: –Identify core and valence electrons –In CHY 116 assign quantum numbers to electrons and more!

36. Terms Electron configuration – shows the number of electrons in each sublevel Box/energy diagram – shows the number of electrons in each orbital –Orbitals are shown as boxes –electrons are shown as arrows

37. Terms Lewis Dot Symbol – shows the valence electrons as dots around the symbol for the element –Maximum of 2 electrons per side of the symbol

38. Terms Valence electrons – all the electrons in the highest occupied PEL –Valence electrons are the ones involved in bonding Core electrons – all electrons not considered valence electrons

39. Still to Come in January Quantum numbers and 7.12, periodic trends