1. Electromagnetic Radiation

2. What is electromagnetic radiation? Electromagnetic radiation (EMR) is a term used to describe all the different kinds of energies released into space by stars such as the Sun.

3. Types of Electromagnetic Radiation Radio Waves TV waves Radar waves Heat (infrared radiation) Light Ultraviolet Light (This is what causes Sunburns) X-rays (Just like the kind you get at the doctor's office) Short waves Microwaves, like in a microwave oven Gamma Rays

4. All these waves do different things. For example, light waves make things visible to the human eye, while heat waves make molecules move and warm up, and x rays can pass through a person and land on film, allowing us to take a picture inside someone's body.

5. Electromagnetic Radiation Light is more than what we can see…

6. Common Properties of EMR They all travel in waves, like the waves at a beach or like sound waves, and also are made of tiny particles. The fact that electromagnetic radiation travels in waves lets us measure the different kinds of radiation by how long the waves are. The wave length of a particular radiation is one way we can tell the kinds of radiation apart from each other.

7. Electromagnetic Radiation Wave Properties of Light: 1. It’s fast! …c = 3.0 x 10 8 m/s 2. It relfects, refracts, diffracts (Transverse wave) 3.

8. All light waves have frequencywavelength symbol: f (Greek “lambda”) units: “cycles per sec” = Hertz“distance” (m, nm) where c = velocity of light = 3.00 x 10 8 m/sec Electromagnetic Radiation c =  f Increasing frequency

9. Electromagnetic Radiation Example: Red light has = 700 nm. Calculate the frequency, f. = 3.00 x 10 8 m/s 7.00 x 10 -7 m  4.29 x 10 14 Hz f = C

10. Electromagnetic Radiation Particle Properties of Light: 1.A particle of light is called a photon 2.Energy of a photon is calculated by E photon = h ·  f where E = energy (Joules, J) f = frequency (Hertz, Hz, 1/sec) h = Planck’s constant 6.63 x 10 J · s -34

11. Electromagnetic Radiation Subatomic particles (electron, photon, proton, etc) exhibit both PARTICLE and WAVE properties. This is known as Wave-Particle Duality. Diffraction: wave-like Photoelectric Effect: particle-like

12. Electromagnetic Radiation Albert Einstein postulates the Photoelectric Effect to explain two observations: 1.No electrons are observed until a minimum energy is applied. 2.Number of electrons ejected depends upon light intensity – not light frequency!

13. Light is created by the Photoelectric Effect

14. Electromagnetic Radiation The photoelectric effect and the idea of discrete, quantized energies neatly explain the observation of emission spectra.

15. Electromagnetic Radiation Example: Red light has = 700 nm. Calculate the energy per photon. E = hf and c = f So f = c/ and E = hc/ E = (6.63 x 10 Js)(3.0 x 10 m/s) 700 x 10 m E = 2.84 x 10 J -34 8 -19 -9

16. Electron Orbitals While thinking about the emission spectrum of hydrogen, Neils Bohr came up with the planetary model of the atom. In this model, electrons can only orbit the nucleus at discrete distances and particular orbital shape. Orbital model of Na Sharp-line spectrum of H Neils Bohr Electron Orbitals

17. Electron Orbitals (n) n = energy level or shell (n = 1, 2, 3, 4, 5, 6, 7) 1. Energy levels are whole numbers 2. The maximum number of electrons in each energy level equals 2n 2. 3. The rows of the periodic table correspond to energy levels. 1. Whole number energy levels – like a standing wave

18. FLAME TEST LAB

19. 3. The rows (periods) of the periodic table correspond to energy levels. Electron Orbitals (n)

20. Electron Orbitals (l) l = subshell (s, p, d, f, g, h, i, j…) 1. s, p, d, and f are named after the four lines in the hydrogen emission spectrum…Sharp, Principle, Diffuse, Fundamental. 2. Each subshell has a different shape 3. The number of subshells in an energy level is equal to the number of the energy level. Energy Level Number of Sublevels Name of sublevels 11s 22s, p 33s, p,d 44s, p, d, f

21. Electron Orbitals (l) 1. Sharp, Principle, Diffuse, and Fundamental refer to the way the spectral lines look. It was thought that electrons traveling between certain energy sublevels produced those certain lines. This was not correct, but the names stuck.

22. Electron Orbitals (l) 2. Each subshell has a different shape s-orbital 1.Has a spherical shape 2.Can hold up to 2 electrons 3.Lowest energy subshell

23. p-orbitals Electron Orbitals (l) 1.Said to have a “dumbbell shape” 2.Can hold up to 6 electrons

25. Electron Orbitals (l) d-orbitals 1.Said to have a “clover leaf” shape 2.Can hold up to 10 electrons

26. combined orbitals d-orbitals

27. Electron Orbitals (l) f-orbitals 1. Can hold up to 14 electrons

28. combined orbitals f-orbitals

29. Electron Orbitals To write a ground-state electron configuration: 1.Determine how many electrons are present. 2.Follow the Aufbau Diagram (Diagonal Rule) Aufbau Diagram

30. Electron Orbitals Example: Write the ground-state electron configuration for nitrogen. 1. Nitrogen has 7 electrons 2. Follow the Aufbau Diagram 3. N: 1s 2 2s 2 2p 3

31. Electron Orbitals So why does it work like this? 1. Pauli Exclusion Principle – states that “no two electrons in an atom can have the same set of four quantum numbers.” In other words, no atomic orbital can contain more than two electrons. 2. Hund’s Rule – The most stable arrangement of electrons around an atom is one with the maximum number of unpaired electrons. This minimizes electron-electron repulsion.

32. Electron Orbitals So why does it work like this? (cont.) 3. Aufbau Principle – Electrons occupy the lowest energy state possible. 4. Heisenberg Uncertainty Principle – The orbitals are probabilities – not shapes in space like planetary orbits. The uncertainty principle states that you cannot know the location and velocity of an electron simultaneously.

33. s-orbitals in Zinc

34. p-orbitals in Zinc

35. Electron Configuration Shortcut…

36. Electron Orbitals Electron orbital notation goes one step further than electron configuration. It describes, specifically, each electron. Compare them Electron Configuration of Oxygen: 1s 2 2s 2 2p 4 Electron Orbital Notation of Oxygen:..... 1s 2s 2p

37. Electron Orbitals OrbitalNotation s. or. or. 1s 2s 3s p... 2p d..... 3d f....... 4f

38. Electron Orbitals Example: What is the electron orbital notation for sulfur?......... 1s 2s 2p 3s 3p Example: What is the non-core electron orbital notation for gold? [Xe]...... 6s 5d

39. Electron Orbitals Example: What is the non-core electron orbital notation for gold? [Xe]...... 6s 5d …or more likely, [Xe]...... 6s 5d Electrons are more stable in full or half- full orbitals.

40. Electron Orbitals Octet Rule: Atoms will gain or lose electrons to achieve a full valence shell (usually this means 8 electrons). Oxidation State: The value of the charge on an ion (positive or negative), after the atom has achieved a full valence shell. - metals tend to lose electrons, forming positive (+) ions (cations). - non-metals tend to gain electrons, forming negative (-) ions (anions).

41. Electron Orbitals Periodic Table of Oxidation States