1. Quantum Theory and the Electronic Structure of Atoms
Chapter 6

2. Objectives 4.0 Define key terms and concepts.
4.1 Demonstrate the relation between wavelength, frequency, and energy.
4.2 Explain the Photoelectric Effect, Shielding Effect of Electrons, and Emission Spectra.
4.3 Explain the dual nature of the electron.
4.4 Explain different models and theories regarding atomic structure.
4.5 Provide the electron configuration, assign quantum numbers to electrons, and diagram the electrons for any element.
4.6 Determine if an element is diamagnetic or paramagnetic.

3. The Nature of Waves Wave Frequency (ν) Wavelength (λ)
An oscillation or vibration in which energy is transferred.
Frequency (ν)
Wavelength (λ)
Amplitude
Wavelength

4. Where c is the speed of light (3.0x108m/s),
v is the frequency,
λ is the wavelength

6. Electromagnetic Radiation
The emission and transmission of energy in the form of electromagnetic waves
Electromagnetic Wave
A wave that has both an electric field and magnetic field with the same frequency and amplitude but vibrate in two perpendicular planes.
X-Rays and Radio Waves
The wavelength is typically measured in nanometers (nm)

8. Planck’s Quantum Theory
The amount of energy emitted by an object depends on the wavelength of the light emitted.
Planck established that atoms can emit or absorb only small quantities of energy.
Quantum
The smallest quantity of energy that can be emitted/absorbed in the form of EM radiation
Planck’s Constant (h) = 6.63x10-34J•s
Where E is energy,
h is Planck’s Constant,
v is the frequency
E=hv

9. The Photoelectric Effect
When electrons are ejected from the surface of a metal as a result of exposure to light above the threshold frequency
Einstein used Planck’s quantum theory to explain the photoelectric effect by stating that light is both wave-like and particle-like
KE=hv-W
Where KE is the kinetic energy of the ejected electron,
W is the work function

10. What is the energy of a photon with a wavelength of 450 nm? For 900nm?

11. What is the wavelength of a photon of light with a frequency of 8
What is the wavelength of a photon of light with a frequency of 8.0x1015/sec? What is its energy in joules?

12. The work function for a metal is 4. 95x10-19J
The work function for a metal is 4.95x10-19J. What is the kinetic energy of the ejected electrons if light with a frequency of 4.15x1015/sec is used to irradiate the metal?

13. The work function for a metal is 8. 15x10-18J
The work function for a metal is 8.15x10-18J. Light with a frequency of 6.00x1015/sec is used to irradiate the metal. How much kinetic energy is produced?

14. What Are Your Questions?

15. Emission Spectra

16. Emission Spectra Emission Spectra Line Spectra
Light emitted by a substance when it is energized using some form of energy
Each element has a unique emission spectra
Line Spectra
Light emitted only at a specific wavelength

17. The Bohr Model of the Atom
Electrons circle the nucleus in different energy levels, each having a specific amount of energy.
The higher the energy level, the higher the energy that electron has and the further the electron is from the nucleus.
Principal quantum number (n)
Positive integers assigned to electron energy levels
The higher the number, the higher the energy

18. The Bohr Model of the Atom
Ground State
The lowest energy level (or shell) an electron can inhabit
Electrons can move up an energy level if it is given enough energy (excited state)
When it returns to its
unexcited stated, a
photon of light (energy)
is released

19. The Dual Nature of the Electron
de Broglie said that if light can act as a particle, then electrons (particles) can act as waves.
The wave-like properties of particles are only observable at the submicroscopic level.

20. Problems with Bohr’s Model
Could not account for the emission of spectra with more than one electron, since his model was based on the hydrogen atom.
Did not explain why additional lines appeared in the emission spectra of the hydrogen atom when a magnetic field was applied.
Cannot identify the location of an electron because they have wave-like properties.

21. Quantum Mechanics Heisenberg Uncertainty Principle
Described the trouble with identifying the velocity and location of a wave
The position or the velocity of an electron can be determined, but not both at the same time.
This indicated that setting electrons in defined orbits around the nucleus was not feasible.

22. Quantum Mechanics Erwin Schrödinger
Used Bohr’s model of the atom and de Broglie’s idea that electrons have wave-like properties to develop the quantum model of the atom, which is still in use today.
Developed Schrödinger’s Equation which mathematically described the behavior and energy of electrons.
Relates the mass and wave-like
behavior of a particle to determine the
likelihood of finding the particle in a
particular area.
Helped start the field of quantum
mechanics.

23. Quantum Model of the Atom
Defines the regions where electrons are most likely to be located at a given time.
Electron Density
High electron density means high probability of finding an electron in that space.
Atomic Orbital
An area where electrons are likely to be
located around the nucleus.
Have a characteristic energy and
electron density.

24. Atomic Orbitals Contain electrons with the same energy
Subshells are identified by letters s, p, d, f
The number of sublevels is equal to the quantum number
Electrons pair up in the sublevels with opposite spins
Each subshell can hold only a certain number of electrons
s can hold up to 2 electrons
p can hold up to 6 electrons
d can hold up to 10 electrons
f can hold up to 14 electrons

25. Atomic Orbitals
p Orbital
s Orbital

26. Atomic Orbitals
d Orbital
f Orbital

27. Electron Configuration

28. Electron Configuration
Aufbau Principle
As protons are added one by one to the nucleus to build up the nucleus of elements, electrons are similarly added to the atomic orbitals.

29. Atomic Orbitals

30. Atomic Orbitals
The higher the Principle Quantum Number, the higher the energy of the electrons held in that shell.

31. Electron Configuration

32. Write the electron configurations for the following elementsNa Ca Fe Br

33. What elements do the following electron configurations represent?
1s22s1
1s22s22p63s23p2
1s22s22p63s23p64s23d104p3
1s22s22p63s23p64s23d104p65s24d5
1s22s22p63s23p64s23d104p65s24d105p4

34. Electron Configurations
Noble Gas Core
Can notate the core electrons using the symbol for the Noble Gas before that element and then write the configuration for the valence electrons.
K is 1s22s22p63s23p64s1 or [Ar] 4s1

35. Write the electron configurations for the following elements using the Noble Gas Core.Na Ca Fe Br

36. What Are Your Questions?

37. Electron Configuration
Orbital Diagram
Show the location and spin of the electrons for an atom
Hund’s Rule
The most stable arrangement of electrons in subshells in the one with the greatest umber of parallel spins.

38. Electron Configuration
Pauli Exclusion Principle
No two electrons can have the same set of quantum numbers
Two electrons occupying the same orbital must have opposing spins
Paramagnetic
Contain unpaired spins and can be attracted by a magnet
Diamagnetic
Does not contain net unpaired spins and is slightly repelled by a magnet.

39. General Rules for Assigning Electrons to Atomic Orbitals
Electrons fill from the lowest shell, moving out from the nucleus.
Each shell (n) contains n number of subshells.
Each subshell of quantum number l contains (2l+1) orbitals.
Cannot pair more than 2 electrons together and electrons will not pair up unless they are forced to.

40. General Rules for Assigning Electrons to Atomic Orbitals
Electrons will not pair up with an electron in the same orbital with the same spin.
The maximum number of electrons in a shell (principal level) is 2n2.

41. Which of the following orbital diagrams are correct?

42. Draw the orbital diagrams for the following compounds
Draw the orbital diagrams for the following compounds. Identify if the element is diamagnetic or paramagnetic.
1s22s22p5
1s22s22p63s23p2 Li B Ca

43. Electron Shielding Effect
As you move out from the nucleus, the orbital shells get bigger.
Electrons held in the outer orbitals will spend less time near the nucleus than those in closer orbitals.
The outer orbitals are shielded by the attractive forces of the nucleus by the inner orbitals.
Electrons in out orbitals require less energy to remove from the atom since they are held less strongly by the nucleus.

44. Quantum Numbers Quantum Numbers Principal Quantum Number (n)
Used to describe the distribution of electrons in atoms
Principal Quantum Number (n)
Whole numbers that indicate orbital level (shell)
Higher the number, the further from the nucleus an electron is located

45. Quantum Numbers Angular Momentum Quantum Number (l)
Indicates the shape of the orbital
Has values from 0 to (n-1), the number corresponds to the subshells present
If n=1, then l=0
If n=2, what is l? l 1 2 3 4 5
Subshell s p d f g h

46. Quantum Numbers Magnetic Quantum Number (ml) Electron Spin (ms)
Describes the location of the subshell in space
For any value of l, there are (2l+1) values for ml, going from –l to +l.
If l=1, then there 3 values for l, -1, 0, and +1
Electron Spin (ms)
Describes the spin on the electron
+½ or -½

47. What is l, ml, and ms if for the second shell? For the fourth shell?

48. Write the quantum numbers for each of the electrons present in Lithium.

49. Write the quantum numbers for each of the electrons present in Nitrogen.

50. Write the quantum numbers for each of the electrons present in Neon.

51. What Are Your Questions?