1. Homework Read pages 360 – 364 372 – 375 379 & 380

2. Electrons in The atom Further developments in the models of the atom

3. Nature Of Light (background Info) Visible light is a form of Electromagnetic radiation Def: form of energy that exhibits wave and particle behavior while traveling through space

4. Wave Characteristics 1. wavelength ( ) – distance between corresponding points on consecutive waves 2. frequency ( f ) – the number of waves that pass through a point in a given amt of time (Hz) 3. Amplitude - the maximum displacement from rest

5. 4. Wave Speed a.Speed =  f b.EM radiation travels at constant speed in a vacuum: c = 3.0 x 10 8 m/s or 186,000 miles/s c.Speed is constant for EM waves, therefore, frequency is inversely proportional to wavelength

6. Frequency vs wavelength

7. EM Spectrum consists of all electromagnetic radiation, arranged according to wavelengths

8. EM Waves and Energy 1.The energy carried by an EM wave can be determined using the following equation: Energy = Planck’s constant x frequency OR E = h f where E = energy in Joules h = 6.63 x 10 -34 J s f = frequency in Hz 2. Energy of the wave is directly proportional to the frequency of that light wave

9. Light from Elements

10. Bohr Model of the Atom The chemical behavior of atoms depends on the arrangement of their electrons -Electrons orbit the nucleus in concentric circular paths He proposed that the reason the electrons ( negatively charged particles) do not fall into the nucleus (positively charged) is because electrons orbit the nucleus with specific fixed amounts of energy

11. Bohr Model of the Atom -Each energy level is associated with a specific amt of energy (called Principle Energy Levels or PEL) -Further from nucleus, greater the energy of an electron -Max # electrons in a PEL = 2n 2

12. Bohr Model (cont’d) Why do atoms give off light? -animationanimation -Electrons occupy the lowest possible energy level (ground state) -If they absorb sufficient energy, they make a quantum leap to higher energy level (excited state)

13. Emission Spectra Excited electrons are unstable and will fall to ground state Give off a burst of energy called a quantum of energy A quantum of energy that falls in visible spectrum called Photon

14. Electron Configurations (Bohr Model) Shows the arrangement of electrons in an atom Each element on your Periodic Table has an electron configuration Indicates the number of electrons in each energy level for that atom

15. Write the electron configuration for an atom of sodium, Na, on the line below. __________________________________ How many electrons occupy the 1 st PEL ? _________ How many electrons occupy the 3 rd PEL? _________ Write a possible electron configuration for an exited atom of sodium. _______________________________

16. Quantum Mechanical Model of Atom Modern model of the atom

17. What was so wrong with the Bohr model of the atom? The Bad could not predict spectra for atoms with more than one electron Heisenberg Uncertainty Principle – the location of an electron cannot be known at any one point in time The Good accurately predicted emission spectra of hydrogen

18. How did it change? Electrons were found to have wave – like properties Scientists began to treat electrons as waves and particles at the same time, developing new ideas on what an atom “looks” like as a result of these properties

19. Quantum Mechanical Model  Electrons are found in areas of definite energy (PELs)  Electrons do not travel in definite paths around nucleus

20. Where are the electrons? Electrons are located in areas of most probable location (orbitals) Visualized as a cloudy like region around the nucleus

21. Orbitals Electrons found in orbitals which are part of a sublevel of each energy level

22. Electron configurations According to the Wave – Mechanical Model

23. Definition Shows the arrangement of electrons in the atom

24. Sublevels of PELS 1.within an energy level, orbitals with different shapes occupy different regions, known as sublevels 2.the # of the principal energy level will identify the possible number of sublevels 3.first 4 assigned are the s, p, d, and f

25. S - sublevel a.s sublevels have the lowest energy b.contains one orbital c.each orbital can hold a max of 2 e - d.has spherical shape

26. P sublevels a.contains 3 orbitals b.Max of 6 e -, along 3 axis c.has peanut shape

27. D Sublevel a.contains 5 orbitals b.has double peanut shape

28. F sublevel a.contain 7 orbitals b.Has most energy of all sublevels

29. Chart of PELs and sublevels PELSublevel# orbitals# electrons 1 2 3 4

30. The Rules for Electron Configurations Aufbau principle – an electron occupies the lowest energy orbital that can receive it Fill order: 1s 2s2p 3s3p3d 4s4p4d4f 5s5p5d5f 6s6p6d6f 7s7p7d7f

31. The Aufbau principle helps us to determine the electron configuration of atoms. Write the electron configuration of an atom of Beryllium (Be) 1. Identify the number of electrons in the atom. Ex) ________

32. 2. Begin to place electrons in the sublevels, by writing the number of electrons that will fit in each sublevel for that atom. Ex) _____________________ 1s 2s2p 3s3p3d 4s4p4d4f 5s5p5d5f 6s6p6d6f 7s7p7d7f

33. Subtract the number of electrons that have been placed in the sublevel from the total number of electrons in the atom (this will tell you how many electrons you have leftover). Continue placing electrons in sublevels, following the fill order, until you run out of electrons for that atom. Ex) _____________________

34. Configurations and the Periodic Table

35. Orbital Notation graphically represents the arrangement of electrons in their energy levels & sublevels Hund’s rule : electrons occupying the same orbital must have opposite spins (we’ll show that with arrows), and electrons will fill one electron per orbital (with identical spin) in a sublevel before they double up.

36. Write the electron configuration for N. 1s 2 2s 2 2p 3 2p______ ______ ______ Increasing 2s______ Energy 1s______

37. Write the electron configuration for Ne 1s 2 2s 2 2p 6 2p______ ______ ______ Increasing 2s______ Energy 1s______

38. Write the electron configuration for O 1s 2 2s 2 2p 4 2p______ ______ ______ Increasing 2s______ Energy 1s______

39. Write the electron configuration for Ti 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 3d______ ______ ______ ______ ______ 4s______ 3p______ ______ ______ 3s______ 2p______ ______ ______ 2s______ 1s______

40. Valence Electrons Definition: electrons that occupy the outermost PEL of an atom -Maximum number of valence electrons is 8 Reason: result of full s and p sublevels - energy levels (clouds) begin to overlap from the 3 rd to the 4 th energy level

41. How many valence electrons are there in the following: Sodium: ________Argon: ________ Oxygen: ________Magnesium: ______ Carbon: ________Strontium: ______

42. Ions Definition: Electrically charged atoms (unequal # of protons and electrons) - formed when atoms lose or gain electrons - in order to have a complete valence energy level (stable configuration) - Possible charges are listed on the Periodic Table

43. Cations Def: a positively charged ion - Formed when atoms lose electrons Ex) Sodium atomSodium ion (Na +1 ) Sodium ion configuration same as Neon

44. Anions Def: negatively charged ions -Formed when atoms gain electrons Ex) Fluorine atomFluoride ion (F -1 ) Fluoride ion configuration same as Neon