Chapter 11

1. Chapter 11 Electrons in Atoms

2. Greek Idea • Democritus and Leucippus • Matter is made up of indivisible particles • Dalton - one type of atom for each element

3. Thomson’s Model • Discovered electrons • Atoms were made of positive stuff • Negative electron floating around • “Plum-Pudding” model

4. Rutherford’s Model • Discovered dense positive piece at the center of the atom • Nucleus • Electrons moved around • Mostly empty space

5. Bohr’s Model • Why don’t the electrons fall into the nucleus? • Move like planets around the sun. • In circular orbits at different levels. • Amounts of energy separate one level from another.

6. Bohr’s Model Nucleus Electron Orbit Energy Levels

7. Bohr’s Model } • Further away from the nucleus means more energy. • There is no “in between” energy • Energy Levels Fifth Fourth Third Increasing energy Second First Nucleus

8. The Quantum Mechanical Model • Energy is quantized. It comes in chunks. • A quanta is the amount of energy needed to move from one energy level to another. • Since the energy of an atom is never “in between” there must be a quantum leap in energy. • Schrodinger derived an equation that described the energy and position of the electrons in an atom

9. The Quantum Mechanical Model • Things that are very small behave differently from things big enough to see. • The quantum mechanical model is a mathematical solution • It is not like anything you can see.

10. The Quantum Mechanical Model • Has energy levels for electrons. • Orbits are not circular. • It can only tell us the probability of finding an electron a certain distance from the nucleus.

11. The Quantum Mechanical Model • The atom is found inside a blurry “electron cloud” • A area where there is a chance of finding an electron. • Draw a line at 90 %

12. Atomic Orbitals • Principal Quantum Number (n) = the energy level of the electron. • Within each energy level the complex math of Schrodinger’s equation describes several shapes. • These are called atomic orbitals • Regions where there is a high probability of finding an electron.

13. S orbitals • 1 s orbital for every energy level • Spherical shaped • Each s orbital can hold 2 electrons • Called the 1s, 2s, 3s, etc.. orbitals.

14. P orbitals • Start at the second energy level • 3 different directions • 3 different shapes • Each can hold 2 electrons

15. P Orbitals

16. D orbitals • Start at the third energy level • 5 different shapes • Each can hold 2 electrons

17. F orbitals • Start at the fourth energy level • Have seven different shapes • 2 electrons per shape

18. F orbitals

19. Summary # of orbitals Max electrons Starts at energy level sublevel s 1 2 1 p 3 6 2 5 10 3 d 7 14 4 f

20. First Energy Level only s orbital only 2 electrons 1s2 Second Energy Level s and p orbitals are available 2 in s, 6 in p 2s22p6 8 total electrons By Energy Level

21. Third energy level s, p, and d orbitals 2 in s, 6 in p, and 10 in d 3s23p63d10 18 total electrons Fourth energy level s,p,d, and f orbitals 2 in s, 6 in p, 10 in d, ahd 14 in f 4s24p64d104f14 32 total electrons By Energy Level

22. Any more than the fourth and not all the orbitals will fill up. You simply run out of electrons The orbitals do not fill up in a neat order. The energy levels overlap Lowest energy fill first. By Energy Level

23. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s

24. Electron Configurations • The way electrons are arranged in atoms. • Aufbau principle- electrons enter the lowest energy first. • This causes difficulties because of the overlap of orbitals of different energies. • Pauli Exclusion Principle- at most 2 electrons per orbital - different spins

25. Electron Configuration • Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to . • Let’s determine the electron configuration for Phosporus • Need to account for 15 electrons

26. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The first to electrons go into the 1s orbital • Notice the opposite spins • only 13 more

27. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 2s orbital • only 11 more

28. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 2p orbital • only 5 more

29. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 3s orbital • only 3 more

30. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The last three electrons go into the 3p orbitals. • They each go into seperate shapes • 3 upaired electrons • 1s22s22p63s23p3

31. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s The easy way to remember • 1s2 • 2 electrons

32. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 • 4 electrons

33. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 2p6 3s2 • 12 electrons

34. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 2p6 3s2 3p6 4s2 • 20 electrons

35. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 • 38 electrons

36. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 • 56 electrons

37. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 • 88 electrons

38. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6 • 108 electrons

39. Exceptions to Electron Configuration

40. Orbitals fill in order • Lowest energy to higher energy. • Adding electrons can change the energy of the orbital. • Half filled orbitals have a lower energy. • Makes them more stable. • Changes the filling order

41. Write these electron configurations • Titanium - 22 electrons • 1s22s22p63s23p64s23d2 • Vanadium - 23 electrons 1s22s22p63s23p64s23d3 • Chromium - 24 electrons • 1s22s22p63s23p64s23d4 is expected • But this is wrong!!

42. Chromium is actually • 1s22s22p63s23p64s13d5 • Why? • This gives us two half filled orbitals. • Slightly lower in energy. • The same principal applies to copper.

43. Copper’s electron configuration • Copper has 29 electrons so we expect • 1s22s22p63s23p64s23d9 • But the actual configuration is • 1s22s22p63s23p64s13d10 • This gives one filled orbital and one half filled orbital. • Remember these exceptions

44. Electron configurations of Ions • Ca+2 has two less electrons than Ca. • Therefore the energy level diagram must accommodate only 18 electrons (isoelectronic to Argon) instead of 20. • 1s22s22p63s23p6

45. High energy Low energy Low Frequency High Frequency X-Rays Radiowaves Microwaves Ultra-violet GammaRays Infrared . Long Wavelength Short Wavelength Visible Light

46. Atomic Spectrum How color tells us about atoms

47. Prism • White light is made up of all the colors of the visible spectrum. • Passing it through a prism separates it.

48. If the light is not white • By heating a gas with electricity we can get it to give off colors. • Passing this light through a prism does something different.

49. Atomic Spectrum • Each element gives off its own characteristic colors. • Can be used to identify the atom. • How we know what stars are made of.

50. These are called discontinuous spectra • Or line spectra • unique to each element. • These are emission spectra • The light is emitted given off.