1. Aim: How did Niels Bohr describe electrons in the atom?
Do Now: Compare and contrast Rutherford/Chadwick’s model of the atom with Niels Bohr’s model of the atom.
In 1912, Bohr took a job in Rutherford’s laboratory in England. Rutherford had just developed a breakthrough model of the atom, but almost immediately, he found a hole in his theory. Since the nucleus is positively-charged and the electrons are spinning around, sort of like the planets around the sun, they would orbit into the nucleus almost instantly. “Atoms can’t exist.” Rutherford was ready to abandon his theory, when Niels Bohr came along and proposed a solution.

2. Bohr wondered… What causes the emission spectra of different elements?

3. “What causes the unique bright line emission spectra for different elements?”1 1 1 1 6 5 4 3 2 2 2 2
purple
blue
green
red

4. Dr. Niels Bohr described the position and behavior of electrons in the atom (1913).
Bohr stated that electrons move in fixed orbits (also known as “principle energy levels” or “shells”) at specific distances from the nucleus.
n=1
n=2
n=3
n=4
n=5
The amount of energy each electron has is determined by which orbit it travels in.
n=6
n=7
Principal Energy Levels (n)
n=1
n=2
n=3
n=4
n=5
Maximum # of Electrons
2e-
8e-
18e-
32e-

5. Draw the Bohr model for an atom of nitrogen in the ground state.
Electron Configuration: 2- 5
Use your glossary to define “ground state.”

6. What can you conclude about all the elements on the periodic table and their electron configurations?
All electron configurations given on the periodic table are in the ground state.
Give the ground state electron configuration for the following elements:
20Ca: ____________
36Kr: ____________

7. Electron Configuration:
Draw the Bohr model for a nitrogen atom in the “excited state” and write the electron configuration.
The amount of energy each electron has is determined by which orbit it travels in.
Excited
Electron Configuration: 2-
4-1

8. Photon
An electron must gain energy (photon) to move to a higher energy level.
1st shell
2nd shell
3rd shell e-
low energy
Nucleus e- 1p 0n
higher energy
Empty space
even higher
energy

9. The excited state is not stable
The excited state is not stable. When an excited electron returns to ground state, it releases excess energy in the form of photons – packets of light energy.
1st shell
2nd shell
3rd shell e-
low energy
Nucleus e- 1p 0n
higher energy
Empty space
even higher
energy

10. Give a possible excited state electron configuration for the following elements:
20Ca: ____________
36Kr: ____________

11. What is an Atomic Emission Spectrum?
The atomic emission spectrum of an atom is a specific pattern of colored lines that is seen when the light emitted from a sample of identical superheated atoms is viewed from a prism.
hydrogen
400 nm
500 nm
700 nm

12. How did Niels Bohr explain Bright Line Emission Spectra?
An electron
jumping from
the 5th shell to
the 2nd shell
An electron
jumping from
the 3rd shell to
the 2nd shell
An electron
jumping from
the 6th shell to
the 2nd shell
An electron
jumping from
the 4th shell to
the 2nd shell
hydrogen
400 nm
500 nm
700 nm

13. What gases comprise the unknown?
Do Now: (Back of notes)
What gases comprise the unknown?

14. Aim: How can we represent the valence electrons of an element?
Do Now: Use your glossary to define the term “valence electrons.”

15. The electrons in the outermost level of an atom.
Do Now
Define “valence electrons.”
What element is this?
How many occupied energy levels?
Identify the number of valence electrons
Name another element with the same # of valence electrons
The electrons in the outermost level of an atom.
Phosphorus
Three 8 5 2
Five
Any element in group 15
(ex: N, P, As, Sb, Bi)

16. Valence Electrons
Valence electrons determine the physical and chemical properties of an atom and are involved in bonding.

17. LEWIS (ELECTRON) DOT DIAGRAMS Illustrates VALENCE ELECTRON CONFIGURATION only.
Write the element’s symbol
Retrieve electron configuration from Periodic Table. The last number in the configuration is the NUMBER OF VALENCE ELECTRONS.
Arrange the valence electrons (DOTS) around the symbol.

18. If you are working with an ion you must adjust the valence electrons (add or subtract electrons) in the configuration before constructing your Dot Diagram – be sure to draw your final diagram with the initial charge on the ion.
Example: Cl-

19. Draw LEWIS ELECTRON-DOT DIAGRAMS for the following:

20. Lab – Identification of Elements
(20 Minutes)

21. Lab – Identification of Elements
(20 Minutes)

22. Lab – Identification of Elements
(20 Minutes)

23. How Neon Lights Work

24. Aim: How are electrons arranged in an atom?

25. Heisenberg Uncertainty Principal (1926)
It is impossible to know both the position of a particle AND its speed/direction at the same time.
In order to see an object, you’re relying on light to bounce off that object and hit your eye. We see electrons in the same way; we try to bounce a photon off of it. Photon hits this electron, electron gains energy and gives off another photon of light. When we detect that photon, however, the electron has now been nudged off of its original course. It’s no longer going in the same speed and direction as it was when the photon hit it.

26. Erwin Schrodinger: The Electron Cloud Model
There is a cloud of probability where you are most likely to find an electron, called “orbitals” (or “electron clouds”).
Basis of the quantum-mechanical model of the atom. The nucleus is not surrounded by “orbits,” it’s surrounded by “orbitals.”

27. 1. Principle Energy Levels
Principle Energy Level (“shell”):
The region where there is the highest probability of locating an electron
Principle Quantum Number (n):
The number given to the shell or principle energy level.
Ex: n=2 refers to the 2nd shell/energy level.

28. Principal Energy Levels correspond to Period #s on your periodic table.

29. 2. Sublevels All principal energy levels have one or more sublevels.
Sublevels are labeled s,p,d, and f.
s<p<d<f describes the order of increasing energy.
The “n” (number of Principle Energy Levels) tells how many sublevels there are.

30. Each Sublevel exists as a different “block” on your periodic table.

31. Principal Energy Level (n)
3. Orbitals
Each sublevel has one or more orbitals. Sublevel s has 1 orbital. Sublevel p has 3 orbitals. Sublevel d has 5 orbitals. Sublevel f has 7 orbitals.
(Each sublevel has two more orbitals than the previous one. Remember the first four odd #s.)
Principal Energy Level (n)
Type(s) of Sublevel
Number of Orbitals 1 s 2 p 3 d 4 f

32. Each Sublevel has one or more orbitals.
s has 1 orbital

33. Each Sublevel has one or more orbitals.
p has 3 orbitals

34. Each Sublevel has one or more orbitals.
d has 5 orbitals

35. Each Sublevel has one or more orbitals.
f has 7 orbitals

36. Principal Energy Level (n) Maximum Number of Electrons
Each orbital can hold a maximum of 2 electrons.
Principal Energy Level (n)
Type(s) of Sublevel
Number of Orbitals
Maximum Number of Electrons 1 s 2 p 3 d 5 4 f 7

37. Sublevels can hold a different number of electrons.

38. The Atom: An Electron Boarding House

39. Electrons always fill the orbitals always fill in a very specific pattern.
You don’t have to memorize these if you understand how the periodic table works.

40. Aufbau Principle: Electrons must fill the lowest energy sublevels before filling higher energy sublevels.
You don’t have to memorize these if you understand how the periodic table works.

41. 8O - _____________________________
Representative Configuration:
Shows how many electrons are in each P.E.L. and sublevel.
Give the extended electron configuration for the following elements:
8O - _____________________________
17Cl - _____________________________
9F - _______________________________
15P - ______________________________
You don’t have to memorize these if you understand how the periodic table works.

42. Noble Gas Configuration:
Shows how many electrons are in each P.E.L. and sublevel in an abbreviated format.
Give the Noble Gas configuration for the following elements:
8O - _____________________________
17Cl - _____________________________
9F - _______________________________
15P - ______________________________
You don’t have to memorize these if you understand how the periodic table works.

43. Rule #2: Pauli Exclusion Principle
An orbital pair of electrons must have opposite spins. One electron spins “up” ( ) and one electron spins “down” ( ).

44. Rule #3: Hund’s Rule
Every orbital in a sublevel must be occupied by a single electron before any of the orbitals can receive a second electron.

45. Remember how many orbitals are in each sublevel!
Orbital Notation:
Uses up and down arrows to show the electron pairs occupying each orbital.
Remember how many orbitals are in each sublevel!
Give the orbital notation for the following elements:
9O - _____________________________
17Cl - _____________________________
9F - _______________________________
15P - ______________________________
You don’t have to memorize these if you understand how the periodic table works.
Remember: s has 1 orbital & p has 3 orbitals.
__ __ __ __ __ __ __ __ __
1s s p s p

46. Aim: How are electrons arranged in an atom?
Period 6 Do Now: Complete the expanded electron configurations for each of the following atoms:
Arsenic
Nickel
Rubidium

47. As 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3 + 4d 5s 4p 3d 4s s p 1 3p 2 d 3 3s
1 2 p 1 3p 2 d 3 3s 4 5 2p 6 7 2s As f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3 1s +

48. Ni 1s2 2s2 2p6 3s2 3p6 4s2 3d8 + 4d 5s 4p 3d 4s s p 1 3p 2 d 3 3s 4 5
1 2 p 1 3p 2 d 3 3s 4 5 2p 6 7 2s Ni f
1s2 2s2 2p6 3s2 3p6 4s2 3d8 1s +

49. Rb 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 + 4d 5s 4p 3d 4s s p 1 3p 2 d
1 2 p 1 3p 2 d 3 3s 4 5 2p 6 7 2s Rb f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 1s +

50. 4d 5s 4p 3d 4s s
1 2 p 1 3p 2 d 3 3s 4 5 2p 6 7 2s Cl f
1s2 2s2 2p6 3s2 3p5 1s +

51. Lab #6 – Electron Configuration
Due on Friday, October 4th.
Do not do “ground vs. excited state questions” yet.