1. Electronic Structure of Atoms
Chapter 10

2. REVIEW: Waves
Electromagnetic radiation is the building block idea of the electronic structure of the atom.
What is wavelength?
How is energy transmitted in the water?

3. REVIEW: Waves What is frequency (f)?
What is the relationship between wavelength and frequency?
As wavelength increase, the frequency decreases

4. REVIEW: Electromagnetic Radiation
What is the speed of light?
All electromagnetic radiation travels at the same velocity, the speed of light.
The wave nature of light does not explain why an object glows when it is hot.

5. Review: Types of Energy
What is kinetic energy?
What is potential energy?

6. The Nature of Energy “Energy” in an atom is a lot like a ladder.
Energy comes in packets called quanta.
Energy levels are fixed energies of an electron.
A quantum of energy is the amount of energy required to move from one energy level to another energy level.

7. Where h is Planck’s constant, 6.626 x 10-34 J-s
The Nature of Energy
Einstein used the assumption of energy levels to explain the Photoelectric effect.
He concluded that energy is directly related (proportional) to the frequency:
E = hv
Where h is Planck’s constant, x J-s

8. Example #1: The Nature of Energy
Make sure you change nm to m
1 nm = 1 E-9 m
Calculate the energy (in joules) of:
(a) a photon with a wavelength of 5.00 x 104 nm (IR region)
(b) a photon with a wavelength of 5.00 x 10-2 nm (X-Ray region)

9. The Nature of Energy
Another mystery in the early twentieth century involved the mission spectra observed from energy emitted by atoms and molecules.

10. The Nature of Energy
For atoms and molecules, one does not observe a continuous spectrum, as one gets from a white light source.
Only a line spectrum of discrete wavelengths in observed.

11. The Nature of Energy
Niels Bohr adopted Planck’s assumption of energy levels and explained the phenomena in this way:
1. Electrons in an atom can only occupy certain orbits (corresponding to certain energies).

12. The nature of energy
2. Electrons in permitted orbits have specific “allowed” energies; these energies will NOT be radiated from the atom.

13. The energy released in in the form of a photon – particle of light
The Nature of Energy
3. Energy is only absorbed or emitted in such a way as to move an electron from one “allowed” energy state to another; the energy is defined by
E=hv
The energy released in in the form of a photon – particle of light

14. The nature of Energy
4. The greater the energy lost by the atom, the greater the energy of the photon

15. The Uncertainty Principle
Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely its position is known.
School hallway analogy
In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself!

16. This is known as Quantum Mechanics.
Erwin Schrodinger developed a mathematical treatment into which both the wave and the particle nature of matter could be incorporated.
This is known as Quantum Mechanics.

17. Quantum Mechanics What are orbitals?
When solving the wave equation, it gives a set of wave function and their corresponding energies.
Each electron has a unique set of quantum numbers that describes its position.
An orbital is described by 4 quantum numbers.

18. The Four Quantum numbers:
Energy Level = n
Type of Orbital = l
*Magnetic quantum number = ml
Spin of the electron = ms
*not going to be studying

19. Quantum Numbers: Energy Levels
Energy levels corresponds to how close the electron is to the nucleus.
The higher the number of ’n’ the greater the distance it is from the nucleus.
When squared, it will also tell you the number of orbitals available to the electron.
Energy Level (n)
Number of Orbitals (n2) 1 2 4 3 9 16

20. Type Of orbital = l There are four types of orbitals: S-orbital  l=0
P-orbital  l=1
D-orbital  l=2
F-orbital  l=3

21. S Orbitals S orbitals are spherical in shape
The radius of the s-orbital increases with the value of n.
Can hold 2 electrons.

22. S Orbitals
Observing a graph of probabilities of finding an electron versus distance from the nucleus, we see that s orbitals possess n-1 nodes, or regions where there is 0 probability of finding an electron.

23. They have two lobes with a node between them.
P Orbitals
They have two lobes with a node between them.
Can hold 6 electrons.

24. D Orbitals
Four of the five d orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center.
Can hold 10 electrons.

25. F orbitals

27. That is, they are degenerate.
Energies of Orbitals
For a one-electron hydrogen atom, orbitals on the same energy level have the same energy.
That is, they are degenerate.

28. Meaning no two electrons were alike.
Energies of Orbitals
As the number of electrons increases, though, so does the repulsion between them.
Therefore, in many-electron atoms, orbitals on the same energy level are no longer degenerate.
Meaning no two electrons were alike.

29. Spin Quantum Number - ms
This idea that no two electrons are the same led to the ms quantum number.
The spin quantum number has only 2 allowed values: +½ and -½

30. Pauli Exclusion Principle
Define Pauli Exclusion Principle:
Therefore, no two electrons in the same atom can have identical sets of quantum number.

31. Hunds’s Rule
Define Hund’s Rule:

32. Aufbau Principle Define Aufbau principle:
*Electrons can be in higher energy levels when they become excited.

33. Writing Electron Configurations
Electron configurations shows the distribution of all the electrons that are in an atom.

34. Orbital Diagrams Be sure to follow the three rules:
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule
Each box represents one orbital.
Half arrows represent the electrons.
The direction of the arrow represents the relative spin of the electron.
Make sure if they are in the same orbital, they have opposite spins!

35. Try this… Draw the orbital diagram for Nitrogen.
Write the electron configuration for Nitrogen.

36. Example Problem #2: Electron Configurations
Write out the electron configuration for Bromine (Br).

37. Does calcium have two more electrons or two less electrons?
Example Problem #3:
Write out the electron configuration for the calcium ion (Ca2+).
Does calcium have two more electrons or two less electrons?

38. Noble gas Electron Configurations
As you start getting higher in atomic number, electron configurations would take forever to write!
Because of this they came up with a way to shorten the configurations.
Look for the noble gas before the element and write it. Then continue with your electron configuration.
For example… let’s look at the element Hassium (Hs #108)

39. Some anomalies
Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.
For example, let’s look at chromium.

40. Some anomalies
This occurs because the 4s and 3d orbitals are very close in energy.
These anomalies occur in f-block atoms, as well.