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. 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 36.0grams? 3. What is the mass of a sample containing 1.204x10 22 atoms.

2. ATOMIC MASS V ATOMIC MASS NUMBER Almost all carbon is one of three isotopes. 1.Write the isotope notation for carbon-14: 2.Write the isotope notation for carbon 12: 3.What is the mass number of carbon 13? 4. Use your periodic table to determine the atomic mass of carbon. Write it here. 5. How many neutrons does carbon have? 6.Learning objective here! Explain the difference between atomic mass and mass number of an element in terms of isotopes, mass, protons, and neutrons.

3. 3.3 Electrons and Light 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. 3.3 Electrons and Light 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: 2.i.* I 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.

5. Quick Review: Models of the Atom Dalton- Indivisible Atom J.J.Thomson discovers subatomic particle “Plum pudding,” model

6. Rutherford ’ s Gold Foil Experiment Alpha particles shot at a thin piece of gold foil did not pass right through with slight deflection. Instead, most passed straight through. Some bounced right back!

7. Rutherford ’ s Gold Foil Experiment Rutherford concluded that Most of the atom is empty space All of the positive charge and almost all of the mass is Concentrated in the tiny core, “ nucleus. ” composed of protons and neutrons. An idea of the size: Atom = football stadium Nucleus = marble

8. 5.3 Physics and the Quantum Mechanical Model Or, “ How do they get all those colors of neon lights? ”

9. Wave Terminology Amplitude = height of wave Wavelength = distance between crests Frequency = number of crests to pass a point per unit of time

10. 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

11. Electromagnetic SpectrumWavelength of Light p140 Visible light is only part of the electromagnetic spectrum:

12. Atomic Spectra When atoms absorb energy, they move to higher energy (excited) levels. When electrons return to the lower energy level, or ground state, they emit light Each energy level produces a certain frequency of light resulting in an emission spectrum

13. 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

14. 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

15. Development of Atomic Models Rutherford ’ s Model l Dense central Nucleus l Electrons orbit like planets l Atom mostly empty space l Does not explain chemical behavior of atoms

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

17. Note Taking Strategies We will Take notes on section 3.3 Electrons and Light ppg92-95 Using an outlining strategy Use phrases in RED as main topic headings. Use phrases in BLUE as subheadings. Important: illustrations are often more important than the text. study each illustration for meaning. sketch the illustration into your notes if it is pertinent Include vocabulary with meaning. If you don ’ t understand a term, don ’ t just copy its definition. READ FOR UNDERSTANDING. WHEN YOU ARE DONE, YOU NEED TO BE ABLE TO EXPLAIN THE CONNECTION BETWEEN LIGHT AND ELECTRONS.

18. Note Taking Strategies We will Take notes on section 3.3 Electrons and Light ppg92-95 Using an outlining strategy Use phrases in RED as main topic headings. Use phrases in BLUE as subheadings. Important: illustrations are often more important than the text. study each illustration for meaning. sketch the illustration into your notes if it is pertinent Include vocabulary with meaning. If you don ’ t understand a term, don ’ t just copy its definition. READ FOR UNDERSTANDING. WHEN YOU ARE DONE, YOU NEED TO BE ABLE TO EXPLAIN THE CONNECTION BETWEEN LIGHT AND ELECTRONS.

19. The Bohr Model l Energy level like rungs of the ladder 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

20. 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!

21. Atomic Emission Spectra Title your paper: Goal: to understand what emission spectra are. We will make a table in which we draw and describe several emission spectra. 1. Put on the diffraction glasses and look outside Find a rainbow. Draw and describe the spectrum you see. SourceDrawingDescription Sunlight Fluorescent bulb

22. Atomic Emission Spectra Title your paper: Goal: to understand what emission spectra are. We will make a table in which we draw and describe several emission spectra. 1. Put on the diffraction glasses and look outside Find a rainbow. Draw and describe the spectrum you see. SourceDrawingDescription Sunlight Fluorescent bulb 2. Repeat with fluorescent lights on in the room. 3. Repeat with neon, helium, and five other elements

23. Quantum Mechanical Model l Energy quantized; 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 1926 Erwin Schrodinger equation described the energy and position of electrons in an atom

24. 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.

25. 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.

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

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

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

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

30. Atomic Orbitals The principal quantum number (energy level) equals the number sublevels

31. 5.2 Electron Arrangement in Atoms Electron Configuration Electrons and nucleus interact to produce most stable arrangement= Lowest energy configuration

32. 3 rules: Aufbau Principle Electrons fill the lowest energy orbitals first Hydrogen has 1 electron 1s 1

33. 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

34. 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

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

36. Practice Problem Electron Configuration for Carbon (atomic number = 6) 1s 2 2s 2 2p 2

37. Practice Problem Electron Configuration for Argon (atomic # = 18) 1s 2 2s 2 2p 6 3s 2 3p 6

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

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

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

41. 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

42. 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

43. 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

44. 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