1. Chemistry Warm Up Some Dimensional Analysis Review.
PLEASE SHOW YOUR WORK USING CONVERSION FACTORS AND DIMENSIONAL ANALYSIS
If 6.02 x 1023 atoms of carbon have a mass of 12.0 grams, what the mass of x 1023 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 102 grams?

2. Chemistry Warm Up: Periodic Table Scavenger Hunt
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
Dalton- Indivisible Atom
J.J.Thomson discovers subatomic particle
“Plum pudding,” model

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

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

7. The Bohr Model 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
A quantum of energy is the amount of energy required to move an electron from one energy level to another

8. 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!

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

10. 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.

11. 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.

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

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

14. •The third energy level has three sublevels, 3s
Atomic Orbitals
•The third energy level has three sublevels, 3s 3p
And 5 3d orbitals py

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

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

17. 5.2 Electron Arrangement in Atoms
Electron Configuration
Electrons and nucleus interact to produce most stable arrangement=
Lowest energy configuration

18. Aufbau Principle Electrons fill the lowest energy orbitals first
3 rules:
Aufbau Principle Electrons fill the lowest energy orbitals first
Hydrogen has 1 electron
1s1

19. 3 rules:
Pauli Exclusion Principal- two electrons per orbital (one spin up, one spin down)
Boron has 5 electrons
1s2
2s2
2p1

20. 3 rules:
Hund’s rule- In orbitals with equal energy levels, arrange spin to maximize electrons with the same spin
1s22s22p3
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.

21. Conceptual Problem p135 1s2 2s2 2p6 3s2 3p3
Electron Configuration for Phosphorus (atomic # = 15)
1s2
2s2
2p6
3s2
3p3

22. Practice Problem 8a p135 1s2 2s2 2p2
Electron Configuration for Carbon (atomic number = 6)
1s2
2s2
2p2

23. Practice Problem 8b p135 1s2 2s2 2p6 3s2 3p6
Electron Configuration for Argon (atomic # = 18)
1s2
2s2
2p6
3s2
3p6

24. Practice Problem 8c p135 1s2 2s2 2p6 3s2 3p6 3d8 4s2
Electron Configuration for Nickel (atomic # = 28)
1s2
2s2
2p6
3s2
3p6
3d8
4s2

25. Practice Problem 9a p135 1s2 2s2 2p1 1
Electron Configuration for Boron (atomic # = 5)
1s2
2s2
2p1
How many unpaired electrons? 1

26. Practice Problem 8c p135 1s2 2s2 2p6 3s2 3p2 2
Electron Configuration for Silicon (atomic # = 14)
1s2
2s2
2p6
3s2
3p2
How many unpaired electrons? 2

27. Exceptions to the Aubau Rule
Copper atomic number=29
1s2
2s2
2p6
3s2
3p6
3d9
4s2
This is the expected electron configuration

28. Exceptions to the Aubau Rule
Copper atomic number=29
1s2
2s2
2p6
3s2
3p6
3d10
4s1
Half-filled and filled sublevels are more stable, even if it means stealing an electron from a nearby sublevel
This is the actual electron configuration.

29. Exceptions to the Aubau Rule
Chromium atomic number=24
1s2
2s2
2p6
3s2
3p6
3d4
4s2
This is the expected electron configuration

30. Exceptions to the Aubau Rule
Chromium atomic number=24
1s2
2s2
2p6
3s2
3p6
3d5
4s1
Half-filled and filled sublevels are more stable, even if it means stealing an electron from a nearby sublevel
This is the actual electron configuration.

31. 5.3 Physics and the Quantum Mechanical Model
Or, “How do they get all those colors of neon lights?”

32. 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

33. 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

34. 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 108
So, as the frequency of light increases, the wavelength decreases

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

36. Wavelength of Light p140
Sample Problem: What is the wavelength of yellow light from a sodium lamp if the frequency is 5.10 x 1014 Hz (Hz = s-1)
Wavelength • frequency = 3.00x108m/s
Wavelength = 3.00x10-8 m/s / frequency
Wavelength = 3.00x108m/s / 5.10x1014 s-1
Wavelenght = 5.88 x 10-7 m

37. Wavelength of Light p140
#14:What is the wavelength of radiation if the frequency is 1.50x1013 Hz (Hz = s-1)?
Is this longer or shorter than the wavelenght of red light?
Wavelength • frequency = 3.00x108m/s
Wavelength = 3.00x108 m/s / frequency
Wavelength = 3.00x108m/s / 1.50x1013 s-1
Wavelength = 2.00 x 10-5 m
Longer than red lightwhich if between 10-6 and 10-7 m

38. 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 = 3.00x108m/s
frequency = 3.00x108 m/s / wavelength
frequency = 3.00x108m/s / 5.00x10-8m
Frequency = 6.00 x 1015 s-1
ultraviolet

39. 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

40. 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

41. 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