1. Chemistry Warm Up Some Dimensional Analysis Review. PLEASE SHOW YOUR WORK USING CONVERSION FACTORS AND DIMENSIONAL ANALYSIS 1.If 6.02 x 10 23 atoms of carbon have a mass of 12.0 grams, what the mass of 1.51 x 10 23 atoms of carbon atoms. Hint: set up the equality that you know. Make two conversion factors and use one to solve the problem. Check your work using dimensional analysis. 2. How many atoms are there in sample of carbon that weighs 30.0grams? 3. How many atoms are there in a sample that weighs 3.60 x 10 2 grams?

2. Chemistry Warm Up: Periodic Table Scavenger Hunt 1.The periodic table is arranged by atomic number, not by atomic mass. Find a sequence of three elements that are arranged by atomic number but not by atomic mass. 2. Find three elements whose symbols don’t seem to have anything to do with their names. Write the name and the symbol for each. 3. There are two rows at the bottom of the periodic table. Use the atomic number to figure out where they fit in to the periodic table. 4. What would the periodic table look like if those two rows were inserted in order of their atomic number? Make a sketch.

3. Chapter5.1 Models of the Atom California State Science Standards Chemistry 1. The periodic table displays the elements in increasing atomic number and shows how periodicity of the physical and chemical properties of the elements relates to atomic structure. As a basis for understanding this concept: g.* Students know how to relate the position of an element in the periodic table to its quantum electron configuration and to its reactivity with other elements in the table. i.* Students know the experimental basis for the development of the quantum theory of atomic structure and the historical importance of the Bohr model of the atom.

4. Chapter5.1 Models of the Atom Objectives Identify the inadequacies in the Rutherford atomic model. Identify the new proposal in the Bohr model.

5. Chapter5.1 Models of the Atom Objectives Describe energies and positions of electrons according to the quantum mechanical model Describe how shapes of orbitals related to different sublevels differ

6. Development of Atomic Models Rutherford’s Model l Discovered dense positive piece at the center of the atom- “nucleus” l Electrons would surround and move around it, like planets around the sun l Atom is mostly empty space l It did not explain the chemical properties of the elements – a better description of the electron behavior was needed

7. The Bohr Model l Electrons orbit the nucleus l Specific circular orbits l Quantum = energy to move from one level to another

8. The Bohr Model l Energy level of an electron analogous to the rungs of a ladder l The electron cannot exist between energy levels, just like you can’t stand between rungs on a ladder l A quantum of energy is the amount of energy required to move an electron from one energy level to another

9. The Bohr Model Energy level of an electron analogous to the rungs of a ladder But, the rungs on this ladder are not evenly spaced!

10. Quantum Mechanical Model l Energy is quantized - It comes in chunks. l A quantum is the amount of energy needed to move from one energy level to another. l Since the energy of an atom is never “in between” there must be a quantum leap in energy. l In 1926, Erwin Schrodinger derived an equation that described the energy and position of the electrons in an atom

11. 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.

12. 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.

13. Atomic Orbitals Schrodinger wave equation gives the probability of finding an electron in a particular location Orbital = the region in which there is a high probability of finding an electron

14. Atomic Orbitals Energy levels (n=1, n=2…) Energy sublevels = different shapes The first energy level has one sublevel: 1s orbital -spherical

15. Atomic Orbitals The second energy level has two sublevels, 2s and 2p There are 3 p-orbitals

16. Atomic Orbitals The third energy level has three sublevels, 3s 3p And 5 3d orbitals pypy

17. Atomic Orbitals The forth energy level has four sublevels, 4s 4p 4d orbitals And seven 4f orbitals

18. Atomic Orbitals The principal quantum number (energy level) equals the number sublevels The number of electrons that can occupy each sublevel equals 2n 2 4p

19. 5.2 Electron Arrangement in Atoms Electron Configuration Electrons and nucleus interact to produce most stable arrangement= Lowest energy configuration

20. 3 rules: Aufbau Principle Electrons fill the lowest energy orbitals first Hydrogen has 1 electron 1s 1

21. 3 rules: Pauli Exclusion Principal- two electrons per orbital (one spin up, one spin down) Boron has 5 electrons 1s 2 2s 2 2p 1

22. 3 rules: Hund’s rule- In orbitals with equal energy levels, arrange spin to maximize electrons with the same spin Nitrogen has 7 electrons Hund’s Rule: Separate the three 2p elecrons into the three available 2p orbitals to maximize the electrons with the same spin. 1s 2 2s 2 2p 3

23. Conceptual Problem p135 Electron Configuration for Phosphorus (atomic # = 15) 1s 2 2s 2 2p 6 3s 2 3p 3

24. Practice Problem 8a p135 Electron Configuration for Carbon (atomic number = 6) 1s 2 2s 2 2p 2

25. Practice Problem 8b p135 Electron Configuration for Argon (atomic # = 18) 1s 2 2s 2 2p 6 3s 2 3p 6

26. Practice Problem 8c p135 Electron Configuration for Nickel (atomic # = 28) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8

27. Practice Problem 9a p135 Electron Configuration for Boron (atomic # = 5) 1s 2 2s 2 2p 1 How many unpair ed electro ns? 1

28. Practice Problem 8c p135 Electron Configuration for Silicon (atomic # = 14) 1s 2 2s 2 2p 6 3s 2 3p 2 How many unpair ed electro ns? 2

29. Exceptions to the Aubau Rule Copper atomic number=29 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 This is the expected electron configura tion

30. Exceptions to the Aubau Rule Copper atomic number=29 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 This is the actual electron configura tion. Half- filled and filled sublevels are more stable, even if it means stealing an electron from a nearby sublevel

31. Exceptions to the Aubau Rule Chromium atomic number=24 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 This is the expected electron configura tion

32. Exceptions to the Aubau Rule Chromium atomic number=24 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 This is the actual electron configura tion. Half- filled and filled sublevels are more stable, even if it means stealing an electron from a nearby sublevel

33. 5.3 Physics and the Quantum Mechanical Model Or, “How do they get all those colors of neon lights?”

34. Goals Describe the relationship between wavelength and frequency of light Identify the source of atomic emission spectra Explain how frequency of emitted light are related to changes in electron energies Distinguish between quantum mechanics and classical mechanics

35. Quick review of wave terminology Amplitude = height of wave Wavelength = distance between crests Frequency = number of crests to pass a point per unit of time

36. Light waves Amplitude = height of wave Wavelength = distance between crests Frequency = number of crests to pass a point per unit of time For light, the product of frequency and wavelength = speed of light, c Frequency Wavelength = 3.00 x 10 8 So, as the frequency of light increases, the wavelength decreases

37. Electromagnetic Spectrum Visible light is only part of the electromagnetic spectrum:

38. Wavelength of Light p140 Wavelength frequency = 3.00x10 8 m/s Wavelength = 3.00x10 -8 m/s / frequency Wavelength = 3.00x10 8 m/s / 5.10x10 14 s -1 Wavelenght = 5.88 x 10 -7 m Sample Problem: What is the wavelength of yellow light from a sodium lamp if the frequency is 5.10 x 10 14 Hz (Hz = s -1 )

39. Wavelength of Light p140 #14:What is the wavelength of radiation if the frequency is 1.50x10 13 Hz (Hz = s -1 ) ? Is this longer or shorter than the wavelenght of red light? Wavelength frequency = 3.00x10 8 m/s Wavelength = 3.00x10 8 m/s / frequency Wavelength = 3.00x10 8 m/s / 1.50x10 13 s -1 Wavelength = 2.00 x 10 -5 m Longer than red light which if between 10 -6 and 10 -7 m

40. Wavelength of Light p140 #15: What is the frequency of radiation if the wavelength is 5.00x10 -8 Hz (Hz = s -1 ) In what range of the electromagnetic specrum is this? Wavelength frequency = 3.00x10 8 m/s frequency = 3.00x10 8 m/s / wavelength frequency = 3.00x10 8 m/s / 5.00x10 -8 m Frequency = 6.00 x 10 15 s -1 ultraviolet

41. Atomic Spectra When atoms absorb energy, Electrons move to higher energy levels. When electrons return to the lower energy level, they emit light Each energy level produces a certain frequency of light resulting in an emission spectrum

42. Atomic Spectra Emission spectra are like a fingerprint of the element We know what stars are made of by comparing their emission spectra to that of elements we find on earth

43. Explanation of Atomic Spectra Emission spectra like a fingerprint of the element We know what stars are made of by comparing their emission spectra to that of elements we find on earth