1. Chemistry Crosby High School
Chapter 5
Atomic Structure
T1 Valence Electrons
T2 Models of Atoms
T3 Electron configuration
Chemistry
Crosby High School

2. Chapter 5 Homework
Read and outline 5.2 (exclude The Heisenberg uncertainty principle and The Schrodinger wave equation) and 5.3

3. Chapter 5 Plan Day Class work 1
Valence electrons and electron dot structures 2
Intro to models; orbital shape 3
Electron configuration & electron lab activity 4
Electron configuration 5
Worksheet 6
Review
HW due 7
Chapter 5 test

4. Valence Electrons The electrons furthest away from the nucleus
Participate in chemical bonds.
Look at the group number to determine an element's number of V.E.
Group Group
Group Group
Group Group
Group Group

5. Practice How many valence electrons do the following elements contain?B Be K Na C F N

6. Electron Dot Diagrams (aka Lewis Dot diagrams)
Show an atoms valence electrons

7. How to draw a Lewis Dot diagram
1. Draw the symbol
2. Determine valence electrons
3. Starting at 12 o’clock and moving clockwise, draw one dot on each side of the symbol, then draw more dots. X

8. Practice as a class
Lewis Dot Diagram for
Carbon
Nitrogen
Potassium

9. Your turn
Draw the electron dot structure for Oxygen, Lithium, and Argon

10. More practice
Page 162, #26-28, 30, 31

11. Atomic Models & Orbital Shape
Key Topics:
Bohr’s model
Quantum mechanical model
Principal energy levels

12. The Atomic Model as of the early 1900s
It was Rutherford’s model
Positively charged core (nucleus) in the middle of empty space. Electrons surrounded the nucleus
Neutrons not discovered until 1932 by Chadwick

13. Revising the Atomic Model
Limitations of Rutherford's Model:
It could not explain the chemical properties of elements.
For example, his model could not explain why metals give off characteristic colors when heated in a flame

14. Neils Bohr (1913) Physicist
Developed a new atomic model: Electrons are found in specific paths, or orbits, around the nucleus
Student of Rutherford’s.

15. Bohr’s Model Each orbit has a fixed energy
Energy increases as you move away from nucleus
Think of it like a ladder: lowest rung has lowest energy; climb the rungs to gain energy. Remember: valence electrons participate in chemical bonds….it is because they have the most energy
Electrons cannot exist between energy levels

16. Energy Levels Fixed energies an electron can have
Quantum = Energy needed to move an electron from one energy level to another

17. Flame Test
Ground state vs. Excited state

18. Schrodinger (1926) Developed the Quantum Mechanical Model
Describes the probability that electrons will be found in certain locations
Most dense where probability is high and least dense where probability is low
Unlike the Bohr model, this model does not specify an exact path the electron takes

19. Concept Check
How can electrons in an atom move from one energy level to another?
By losing or gaining an amount of energy – a quantum

20. Youtube Video
(6 minutes of video)

21. Review Thomson Rutherford Bohr Schrodinger
Subatomic Particles in the model
Draw a picture of the model
Write a 1 sentence summary of the model
Thomson
Rutherford
Bohr
Schrodinger

22. Quantum Mechanical Model and Energy Levels
Instead of restricting electrons to limited circular orbits, Schrodinger developed a new idea of orbitals
Orbitals describe the probability of finding an electron at various locations around the nucleus.
Schrodinger’s numbers: size, shape and orientation. We will only cover size and shape

23. Probability of Finding an electron
The region of space in which the electron can probably be found is called an orbital.
In this lab, you will use a felt-tip maker and a target to investigate the probability distribution of marks about a central point.

24. What you will do:
Place the target on a notebook or a stack of papers to create a cushioned landing.
2. Drop the marker 100 times from hip-height
3. Create a chart and record the number of marks in each region of the target.
4. Create a graph of your marks. Draw a smooth curve that shows the general trend. DO NOT simply connect the dots.
5. Write a conclusion that explains how the probability of finding an electron changes as you move outward from the nucleus.

25. Electron Arrangement in Atoms (In the Q.M.M.)
Main energy levels (aka: quantum numbers)
- n = 1, 2, 3, 4
Sublevels
- s, p, d, f
The sublevels contain the orbitals

26. Sublevels Each sublevel contains orbitals
Orbitals have specific shapes
s: 1 orbital – sphere
p: 3 orbitals – dumbells
d: 5 orbitals – pears
Look on page 131

27. Summary of Principal Energy Levels and Sublevels (p. 155)
Number of sublevels
Type of sublevel
Maximum # of Electrons
n = 1 1
1s (1 orbital) 2
n = 2
2s (1 orbital), 2p (3 orbitals) 8
n = 3 3
3s (1 orbital), 3p (3 orbitals),
3d (5 orbitals) 18
n = 4 4
4s (1 orbital), 4p (3 orbitals),
4d (5 orbitals), 4f (7 orbitals) 32

28. Review
Page 167, #59-61, 63, 64, 68, 70,
A4 finished #1

29. Electron Configuration
Key Question:
What are the rules for writing the electron configurations of elements?

31. Electron Configurations
Definition: the ways electrons are arranged in various orbitals around the nuclei of atoms.
Three rules tell you how to find the electron configuration of an atom:
Aufbau Principle
Pauli Exclusion Principle
Hund’s rule

32. The Rules 1. Aufbau Principle
electrons occupy orbitals of the lowest energy first
2. Pauli Exclusion Principle
Each orbital holds only 2 electrons
remember s has 1 orbital, p has 3 orbitals and d has 5 orbitals
So, how many electrons will you find in
s sublevel?
p sublevel?
d sublevel?

33. What does all that mean? When writing electron configurations
1.) Follow this chart (Aufbau Principle) (5-C)
Chart is in lab manual

34. 2) Remember: ‘s’ orbital holds a max of 2 electrons, ‘p’ orbital holds a max of 6 electrons and ‘d’ orbital holds a max of 10 electrons
3) Remember, the atomic number tells you how many electrons are in each atom

35. Let’s Practice
Phosphorous has an atomic number of 15, what is the electron configuration?
Follow the chart
1s2
2s2
2p6
3s2
3p3
= 1s22s22p63s23p3

36. More Practice Electron configuration for Carbon (atomic number = 6):
1s22s22p2
Electron configuration for Argon (atomic number = 18)
1s22s22p63s23p6

37. Individual Practice On page 136, answer #8c and all of 9.
- For # 9, just write the electron configuration, do not worry about unpaired electrons.

38. More Practice
In your lab manual, complete pages 5-C (front and back) and the front of 5-D

39. Drawing atomic orbitals:
Use the Aufbau chart to figure out electron configuration
2) Draw arrows to represent electrons
**Pauli exclusion: Electrons spin in opposite directions**
3) Draw all ‘up’ arrows to represent electrons before drawing any ‘down’ arrows
**Hund’s rule: Every orbital in a sublevel is occupied by one electron before any one sublevel is occupied by two electrons**

40. Look at Table 5.2 Remember: s has 1 orbital p has 3 orbitals
d has 5 orbitals
***You must memorize the number of orbitals in each sublevel.

41. Practice
Here is the electron configuration for Phosphorus: 1s22s22p63s23p3
Draw the electron orbital configuration
1s s p s p
__ __ __ __ __ __ __ __ __

42. More Practice On the back of 5-D in your lab manual try # 1 and 2
Do 1 and 2 as a group, if they get those right, they may continue. A4 stopped here

43. Practice
1. Page 136, #9
Draw the electron orbitals for the each atom to determine the number of unpaired electrons
2. On page 137, answer questions 11 and 14

44. Chapter Review
Page 152, # 35-44, 56, 59, 60, 69, 73, 77
Page 157, # 1-3,