2. Something Smaller Than An Atom? Atomic Structure

3. Review of the Atom Smallest part of matter representing an element Once was thought to be the smallest part of matter Later, scientists discovered atoms are made of subatomic particles

4. Subatomic particles Protons - positive charge Neutrons – no charge (neutral) Electrons – negative charge

5. Representing a Specific Element Mass Number (P + + N o ) Atomic Number (P + ) AXZAXZ Atomic Symbol

6. Nucleus: Center Stage Ernest Rutherford discovered atoms have a nucleus (1911) Later scientists discovered that the nucleus also contains all the neutrons

7. Rutherford’s Model Discovered the nucleus Electrons moved around in Electron cloud Mostly empty space

8. Where are those electrons? Early models of the atom showed electrons spinning around the nucleus randomly Research showed that this is NOT true

9. Bohr’s Model Why don’t the electrons fall into the nucleus? Move like planets around the sun. In circular orbits at different levels. Energy separates one level from another.

10. Bohr’s Model Nucleus Electron Orbit Energy Levels

11. More about Bohr….. Created a model showing electrons are in orbits of different energy around the nucleus …… Think of the planets orbiting the sun Bohr used the term energy levels (or shells) to describe these orbits with different amounts of energy. He also said that the energy of an electron is quantized, meaning electrons can have one energy level or another but nothing in between.

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

13. Let’s get close to the nucleus  Bohr found that the closer an electron is to the nucleus, the less energy it needs.  The farther away it is, the more energy it needs  He numbered the electron’s energy levels

14. Energy levels 1 st level can hold up to 2 electrons 2 nd level can hold up to 8 electrons 3 rd level can hold up to 18 electrons And so on……..

15. Energy levels E-level an electron normally occupies is called ground state But, electrons can move to a higher – energy, less-stable level, or shell, by absorbing energy. This higher energy, less-stable state is called the electron’s excited state.

16. The electron gets crunk….. When the electron is finished being excited it goes back to its ground state by releasing some of the energy it has absorbed

17. Line spectrum…..what is that??? Energy released by electrons sometimes occupies part of the electromagnetic spectrum that humans detect as visible light

18. Problem with Bohr’s model Unexplainable observations on complex atoms until the quantum theory was created Quantum theory ----- matter has properties associated with waves. It is impossible to know the exact position and momentum (speed and direction) of an electron at the same time. ------UNCERTAINTY PRINCIPLE

19. The Quantum Mechanical Model Energy is quantized. It comes in chunks. Quanta - the amount of energy needed to move from one energy level to another. Quantum leap in energy. Treated electrons as waves

20. Does have 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. The Quantum Mechanical Model

21. The electron is found inside a blurry “electron cloud” A area where there is a chance of finding an electron. The Quantum Mechanical Model

22. Quantum mechanical model of the atom Pay attention so you don’t get LOST !!!!!!!

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

24. 1 s orbital for every energy level Spherical shaped Each s orbital can hold 2 electrons Called the 1s, 2s, 3s, etc.. orbitals. S orbitals

25. P orbitals Start at the second energy level 3 different directions 3 different shapes (dumbell) Each can hold 2 electrons

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

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

28. Summary s p d f # of shapes Max electrons Starts at energy level 121 362 5103 7144

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

30. By Energy Level Third energy level s, p, and d orbitals 2 in s, 6 in p, and 10 in d 3s 2 3p 6 3d 10 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 4s 2 4p 6 4d 10 4f 14 32 total electrons

31. Bed check for Electrons…Electron Configurations…… Goal: Use an energy level diagram to depict electrons for any element 1s orbital is closest to the nucleus and it has the lowest energy At energy level 2, there are both s and p orbitals, with the 2s having lower energy than the 2p.

32. Energy level diagram continued The three 2p subshells are represented by three dashes of the same energy. Energy levels 3, 4, and 5 are also shown Notice, that 4s has lower energy than the 3d. This is an exception to what you thought but it does occur in nature like this.

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

34. 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

35. Aufbau Principle Method for remembering the order in which orbitals fill the vacant energy levels (like a people fill up vacant rooms in a hotel) RULE: ELECTRONS FILL THE LOWEST VACANY ENERGY LEVELS FIRST. ANOTHER RULE: WHEN THERE’S MORE THAN ONE SUB-SHELL AT A PARTICULAR ENERGY LEVEL, SUCH AS AT THE 3P OR 4 D LEVELS, ONLY ONE ELECTRON FILLS EACH SUB-SHELL UNTIL EACH SUBSHELL HAS ONE ELECTRON. THEN, ELECTRONS START PAIRING UP IN EACH SUBSHELL. THIS RULE IS CALLED HUND’S RULE.

36. 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 Phosphorus Need to account for 15 electrons

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

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

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

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

41. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f The last three electrons go into the 3p orbitals. They each go into separate shapes 3 unpaired electrons 1s 2 2s 2 2p 6 3s 2 3p 3

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

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

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

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

46. Draw the energy level diagram of oxygen Find oxygen on the period table Atomic number = 8 Therefore, it has 8 protons and 8 electrons So, you put 8 electrons into your energy level diagrams

47. Energy level diagram for oxygen Use arrows to represent electrons If two electrons fit in the same orbital, one must face up and the other one down The first electron goes into the 1s orbital, filling the lowest energy level first. And, the second one spin pairs with the first one. Electrons 3 and 4 spin pair in the next lowest vacant orbital – the 2 s. Electron 5 goes into one of the 2p sub-shells Electrons 6 and 7 go into the other two totally vacant 2p orbitals The last electron spin pairs with one of the electrons in the 2p subshells.

48. Electron Configuration Oxygen 1s 2 2s 2 2p 4 You can derive the electron configuration from the energy level diagram. The first two electrons in oxygen fill the 1s orbital, so you it as 1s 2 in the electron configuration. The 1 is the energy level, the s represents the type of orbital, and the superscript 2 represents the number of electrons in that orbital. The next two electrons are in the 2s orbital, so you write 2s2. Last, you show the 4 electrons in the 2p orbital as 2p4. That’s how you get 1s 2 2s 2 2p 4.

49. Tip The sum of the superscript numbers equals the atomic number, or the number of electrons in the atom.

50. Exceptions to Electron Configuration

51. Orbitals fill in order Lowest energy to higher energy. Adding electrons can change the energy of the orbital. Filled and half-filled orbitals have a lower energy. Makes them more stable. Changes the filling order of d orbitals

52. Copper’s electron configuration Copper has 29 electrons so we expect 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 But the actual configuration is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 This gives one filled orbital and one half filled orbital. Remember these exceptions  s 2 d 4  s 1 d 5  s 2 d 9  s 1 d 10

53. Here are a couple of electron configurations you can use to check your conversions from energy level diagrams: Chlorine (Cl) Iron (Fe)

54. Answers Answer: Cl 1s 2 2s 2 2p 6 3s 2 3p 5 Answer: Fe 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6