1. Chapter 6 – Electronic Structure of Atoms

2. Electromagnetic Spectrum (Light)

3. We can use wavelength and frequency to calculate the energy in a photon (particle of light).
low frequency
high frequency
Wavelength (l) is the distance between two crests on a wave.
Frequency (f) is the number of waves that pass through a particular point in 1 second (Hz = 1 cycle/s).

4. What is the relationship between energy and frequency?
Energy Increases
Frequency Increases
Energy and frequency are directly related
If frequency increases then energy increases
If energy decreases then frequency decreases

5. What is the relationship between energy and wavelength?
Energy Increases
Wavelength Decreases
Energy and wavelength are inversely related
If wavelength increases then energy decreases
If wavelength decreases then energy increases

6. What is the relationship between frequency and wavelength?
Frequency Increases
Wavelength Decreases
Energy and wavelength are inversely related
If frequency increases then wavelength decreases
If frequency decreases then wavelength increases

7. Why do we care about light in Chemistry?!
Electrons release light when they are excited and then calm back down. e-
Light (Photon – particle of light)
Examples of light are gamma rays, x-rays, radio waves, microwaves
The particle of light released is called a photon

8. Properties of Light
All electromagnetic radiation travels at the same velocity: Namely the speed of light in a vacuum (c) is 3.00  108 m/s.
c = 
Ex: A certain AM radio station broadcasts at a frequency of 6.00x102 kHz. What is the wavelength?

9. Properties of Light
Light can be defined as both a PARTICLE and a WAVE.

10. Properties of Light
Energy (light) is emitted or absorbed in discrete units (quantum)
Each metal has a different energy at when it emits electrons. At lower energy, electrons are not emitted.
Einstein used quanta to explain the photoelectric effect. Energy is proportional to frequency.
E = h
where h is Planck’s constant, h = 6.63  10−34 J∙s.

11. Properties of Light
Ex: A given light has energy of 2.85x10-19 J. Calculate the frequency and wavelength of that light.

12. The Bohr Model ( ) E = RH –
Electrons can have specific (quantized) energy values.
Light is emitted as e- moves from one energy to a lower energy level
Energy is depicted in the following equation:
where RH is the Rydberg constant, 2.18  J, and ni and nf are the initial and final energy levels of the electron.
E = RH
( ) 1
ni2
nf2 –

13. Energy Practice
Ex: A photon has absorbed energy when it jumped from n = 2 to n = 4 level. Calculate:
Energy
Frequency
Wavelength of the photon

14. Electron Configurations
As chemists, we need to know where an electron is at so we can manipulate it and design awesome stuff
Image of iron atoms
on copper surface

15. Moving away from the Bohr Model…
If negative electrons actually orbited a positive nucleus, they would eventually spiral in and collide, causing an explosion.

16. In reality, the position of an electron is a matter of probability.
The space where an electron will probably be found is called an “orbital”
Probably here
Probably NOT here

17. Quantum Numbers
Four quantum numbers are required to describe the distribution of electrons in atoms.
Each orbital describes a spatial distribution of electron density.

18. Principal Quantum Number (n)
The principal quantum number, n, describes the energy level of the orbital
The values of n = 1,2,3,4…
It gives you the distance of e- from the nucleus

19. Angular Momentum Quantum Number (l)
This quantum number gives the shape of the “volume” of space that the e- occupies
l = 0,1,2,3 … n-1

20. Magnetic Quantum Number (ml)
The magnetic quantum number describes the orientation of the orbital in space.
ml = −l, … 0, … +l
If l = 1 (p orbital), then ml = -1, 0, or 1
If l = 2 (d orbital), then ml = -2,-1, 0, 1, or 2
If l = 3 (f orbital), then ml = -3,-2,-1,0,1, 2, or 3

21. Relation to Quantum Numbers and Orbitals

22. Magnetic Spin Quantum Number (ms)
Spin quantum number displays the orientation of the electron within the orbital
There can only be 2 electrons in an orbital, where the ms value is either -1/2 or +1/2
+1/2
-1/2

23. Subshell Shapes
These shapes represent where the electrons are probably hanging out

24. The Four Subshells s Subshell Shape Max # of electrons # of orbitals
Sphere 2 1

25. The Four Subshells p Subshell Shape Max # of electrons # of orbitals
Dumbbell 6 3

26. The Four Subshells d Subshell Shape Max # of electrons # of orbitals
4-Lobed 10 5

27. The Four Subshells f Subshell Shape Max # of electrons # of orbitals
6-8 Lobed 14 7

28. Now for a couple rules:
Aufbau Principle: add electrons from lowest energy to highest energy (you pretty much already do this when you use the proper filling order)
Pauli Exclusion Principle: There is a maximum of two electrons for each line (or orbital)
Hund’s Rule: Electrons fill in each orbital first before they start pairing up with an electron

29. Using the periodic table makes it sooooo easy
Using the periodic table makes it sooooo easy! All you have to do is memorize this:

30. This is what the periodic table would
look like if we had more room

31. Or…you can memorize this:

32. What is an electron configuration?
It identifies what and where we can expect to find an atom’s electrons
The electron configuration for Hydrogen
1s1
Number of e- in subshell
s: max of 2
p: max of 6
d: max of 10
f: max of 14
Shell (could be 1-7)
Subshell (s,p,d,f)

33. Warm Up
1. What are the 3 orbitals and how many electrons can each orbital hold? 2. What is the next orbital in the series of filling up of the electrons? 1s, 2s, 2p, 3s,3p,__ 3. Given the electron configuration of 1s2 2s2 2p5 What is the element?

34. Writing the e- configuration:
Steps:
Count the number of electrons in the element
Use the periodic table to follow the order of filling
As you start writing keep count of the electrons used
Remember: s [2e-] p [6e-] d [10e-] f [14e-]
Check the number of electrons when finished
Examples:
1) Carbon [6e-]
2) Sodium [11e-]
1s2
2s2
2p2
1s2
2s2
2p6
3s1

35. Use the nearest noble gas as a short cut. [He, Ne, Ar, Kr, Xe…]

36. Longhand Configuration Shorthand Configuration
Sulfur 16e-
1s2
2s2
2p6
3s2
3p4
Shorthand Configuration
Sulfur 16e- [Ne] 3s2 3p4

37. Orbital Diagrams Each box in the diagram represents one orbital.
The arrows represent the electrons, where each arrow indicates the spin