Electrons in the Atom

1. Electrons in the Atom Part 2

2. Remember Atomic Theories� What was wrong with Rutherford�s model of the atom? Electrons were outside the nucleus and need some method of organization What �exciting� scientist came up with a theory to organize electrons? Niels Bohr

3. Niels Bohr Bohr Model Electrons can only possess certain amount of energy Electrons must be in specific energy level Need more energy to be farther away from the nucleus Electrons CANNOT be in between energy levels

4. Bohr Model of Atom

5. Bohr Model

6. Hydrogen- emitting a photon

7. Electron Energy Levels Energy Levels = the possible electrons orbits of an atom Ground State = exists when an atom is energetically stable Excited State = exists when electrons absorb energy, move to higher energy levels, atom becomes energetically unstable.

8. Vocabulary When an atom is put into a flame describe what happen to electrons using these vocabulary words: energy levels, ground state, excited state

9. Quantum Mechanical Model Currently accepted model energy is quantized energy has only certain allowable values; other values are NOT allowed Cannot tell exact location of electron in quantum mechanics Can give the probability of finding any electron a certain distance from the nucleus. Can identify the energy an electron is in

10. Electrons & Light Continuous Spectrum = band of colors that results when a narrow beam of light passes through a prism Many electrons jump energy levels all at once, so lots of different colors are emitted

11. Bright Line Spectrums Each element has its own set of lines Electrons drop from different levels in different atoms Unique like finger prints.

12. Emission Spectra- Finger Prints

13. Electron Motion for Hydrogen

14. Electrons in Atom There are 7 energy levels in an atom Higher Energy Level means� Farther the electron is from the nucleus The more energy the electron has

15. Energy Levels Sublevels (Area within energy level) Each have different energy value Called Orbitals (like apartments on a floor) s, p, d, f Each orbital can hold a max of 2 electrons

16. Orbitals

17. Orbitals & Energy Levels (n) n = energy level in the atom n = 1st = s n = 2nd = s, p n = 3rd = s, p, d n = 4th = s, p, d, f n = 5th = s, p, d, f n = 6th = s, p, d,

18. Orbital Shapes

19. Order to fill Orbitals Low energy first, then higher energy

20. Electron Configuration Shows arrangement of all the electrons Aufbau Principle = fill lowest E orbital to highest E orbital. Paulii- Exclusion Principle = max of 2 electrons in each orbital; electrons have opposite spins (think ying yang) Hund�s Rule = each same spin electron must fill an empty orbital of equal energy before doubling

21. Electron Configuration Format : 1s1 1 = the energy level s = the sublevel, or orbital 1 = the number of electrons in that sublevel Written as an exponent (superscript)

22. Order to Fill Orbitals

23. Fill According to Periodic Table

24. Writing Electron Configuration Locate the element on the periodic table Fill in the orbitals in the proper order Check that total number of electrons equals the atomic number

25. Electron Configuration

26. Examples C = Li = Na = Cl = K = Fe =

27. Shorthand Notation Put noble gases that precedes the element in brackets, then continue filling the rest in order. Na = Cl = K = Fe =

28. Orbital Notation Same as electron configuration, use arrows to represent electrons Lines or circles to represent orbitals Example- Fluorine

29. Significance of E- Configuration Valence Shell = the outermost energy level of an atom Valence electrons = electrons in the valence (outer) shell; important because involved in chemical bonding Max number of valence electrons an atom can have is 8 s and p orbital (2 + 6 = 8 electrons)

30. Significance of Noble Gases Noble Gases = not reactive Full outer shell, s and p orbitals Every other element wants to have an electron configuration like a Noble Gas