1. Quantum Mechanics Periodic Trends Chemical Bonding
Electrons in Atoms
So why does potassium explode in water?
Potassium metal explosion
Quantum Mechanics
Periodic Trends
Chemical Bonding

2. 12.1 Development of Atomic Models
Thompson’s
Dalton’s
Rutherford’s
Quantum Model
of helium shows
2 electrons in the
1s orbital or “1s2”
Bohr’s carbon
Bohr proposed electrons orbit in paths of fixed energy called energy levels.

3. Size of Atom and Subatomic particles

4. Quantum Theory Symphony

5. Quantum Mechanical Model
The quantum model of the atom is based on the solution to the Schrödenger equation.
One way to visualize the model’s electron levels is to imagine a ladder where the higher rungs or levels are closer together.
7 rungs = 7 energy levels
Principle quantum numbers (n) correspond to the energy levels

6. Atomic Sublevels or S and P orbitals
Probability cloud models (left) show where it is most likely to find electrons. For each principal quantum number (n) there are the same number of sub levels. (n=2, 2 sublevels)
Helium n=1, had 1 suborbital the “s” orbital. Level 2, n=2, has 2 sub levels, “s” and “p.”
(The px,py and pz orbitals are found on energy levels 2-7)
Each energy level, like level 2 shown here in green. has a spherical “s” orbital.

7. How many sublevels are there for each principle quantum number?
Ask Your Neighbor:
How many sublevels are there for each principle quantum number?
Which energy levels contain px, py and pz orbitals?
How are “s” orbitals different than “p” orbitals?

8. Electron Configuration of Hydrogen4s
3 d
How are all group IA elements similar in electron configuration? 3p
Increasing energy 3s 2s
2px 2py 2pz
H 1s1 1s

9. Light will bend and reflect at the interfaces between different materials.
Prism lenses bend light. White light is a blend of all wavelengths of visible light.

10. Each color has a different wavelength (λ) lambda = wavelength
1000 µm = 1 mm

11. First time for everything: Light as a particle and a wave!
Light photographed as a wave and as a particle first time!

12. The shortest wavelength also has the highest energy, hence UV light can harm us if the wavelength is too short!

13. Evidence of Energy Levels and Suborbitals
What are emission spectra?
How is each element’s emission spectra unique?

14. Discrete lines = quanta of energy
Why does each element have its own “signature” emission spectrum? Tell neighbor.
A. Each element has a different number of protons and electrons.
B. Each element has unique nuclear attraction for electrons in shells.
C. Each atom’s first energy level is a unique distance from nucleus.
D. Distance between outer energy levels in atom is unique to each element.
E. Electrons emit photons whose frequency is proportional to energy lost.

15. Predictions Based on Models of the Atom
Click here

17. Electron Configuration Rules
Aufbau principle- Electrons enter orbitals of lowest energy first.
For order of orbitals from lowest to highest learn the ZIG ZAG RULE.

18. Zig-Zag Rule #e- Atomic # 8 92 118 at 7p filled 18 60 zig 6 7s 7p
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p 1s

19. Electron Configuration of Helium4s
3 d 3p
Increasing energy 3s 2s
2px 2py 2pz
He 1s2
1ss
(filled)

20. Pauli exclusion principle
Hund’s Rule
Pauli exclusion principle
In a single energy level, electrons must occupy only one orbital until each orbital has an electron.
Since electrons have the same charge, they have strong repulsion forces, pushing them into different suborbitals of the same energy.
Orbits fill with electrons of opposite spin. (½+
and ½-)
Electrons (e-) have strong, negative repulsion forces, but they are less repulsive if they have opposite spin.

21. Electron Configuration of Nitrogen4s
3 d 3p
Increasing energy
Predict the electron configuration of Ne. 3s 2s
2px 2py 2pz
N 1s2 2s2 2p3
1ss

22. Electron Configuration of Neon4s
3 d
Predict the electron configuration of Ar. 3p
Increasing energy 3s
Filled 2nd energy level
(8 electrons = octet) 2s
2px 2py 2pz
Ne 1s2 2s2 2p6
1ss

23. Electron Configuration of Argon4s
3 d
3rd energy level
(8 electrons = octet) 3p
Increasing energy 3s
Filled 2nd energy level
(8 electrons = octet) 2s
2px 2py 2pz
Ar 1s2 2s2 2p6 3s2 3p6
1ss

24. Observe this Shockwave Electron Configuration which introduces us to the mystery of some periodic trends.
1. What happens to the size of the atom as the energy levels are filled?
2. When do the energy levels change the most?

25. High frequency light frees electrons
Photoelectric Effect
Click here for photoelectric effect simulation
High frequency light frees electrons
from reactive metals

26. Periodic means cycle or repeating pattern.
Periodic Trends
Patterns in the physical and chemical properties of the elements are called trends.
Periodic means cycle or repeating pattern.

27. Periodic Trends in Atomic radius
Group trends- Radius increases as electrons (must fill new energy levels) and are added to atom.
Atomic mass, and number increase in this direction also. r

28. What happened to energy levels as p+ and e- increased across a row?
How does this affect atomic radius across a row?

29. Periodic Trends in Atomic radius
Period trends- Radius decreases as electrons fill across same energy level. Filled inner levels shield outermost electrons from the nucleus
(So in any period, between the nucleus and outer electrons, there is the same number of electrons.)
This trend is opposite for atomic mass & number.

31. Electronegativity Metal vs. Nonmetal
Where are the metals vs. the nonmetals?
Nonmetals have high electronegativity.
Metals have low electronegativity

32. Electronegativity vs. atomic size
Where are the smallest atoms in a period?
Big atoms have lower electronegativity

33. Electronegativity explained
Valence electrons of small atoms that are closer to the nucleus than larger atoms, tend to be held to the nucleus with stronger forces of attraction. Usually the farther they are away, the weaker the forces of attraction. e- e-
High or strong attraction to valence electrons in a bond = High Electronegativity.

34. First Ionization Energy
This is the amount of energy it takes to remove the first or outer most electron.
Look on your periodic table at first ionization potential in V, or on page in textbook.
How easy is it to remove electron from the Group I & II metals? From the halogens?

35. First Ionization Energy

36. How might the first ionization energy compare to the electronegativity across the first period?

37. They are very similar!

38. Trends Important to Bonding
Ionization energy was used to help determine electronegativity.
Electronegativity is a scale in the units of Paulings, developed or calculated to show the degree one element tends to have the bonding electron(s) in a pair of oppositely charged ions.
NaCl e- e-

39. For example + Cl
3.12 K
0.82 -
The electronegativity of nonmetal Cl- is 3.12
And the electronegativity of metal K+ is .82
The difference between the two is 2.30 Pauling units
We determine the percent ionic character of the bond to be 74%. By definition this is considered ionic bonding since it is more than a 2.0 difference.

40. To determine bond type: Calculate the difference in Electronegativity between these Element Pairs
Difference > 2.0 = ionic
Difference < 2.0 = covalent
0.35
3.0
C and H
Li and F
0.94
Na and F
3.05
N and H
K and F
3.16
0.48
S and H
(Formula Units of bonded ions)
1.43
C and F
(Molecules of bonded atoms)

41. Write electron configurations of the following:
1. Ca Br-1
2. Al S-2
3. K N
4. Kr [Ar] 4s2 3d10 4p6
5. Ar 1s2 2s2 2p6 3s2 3p6
6. 1s2 2s2 2p3
1. Ar 1s2 2s2 2p6 3s2 3p6
2. Ne 1s2 2s2 2p6
3. 1s2 2s2 2p6 3s2 3p64s1

42. Use electron dot diagrams to determine chemical formulas of the ionic compounds formed when the following elements combine.
Example:
K and I
Ca and S +1 -1 -2 +2 K I Ca S