1. Electrons in atoms
Chapter 5

2. 5.1Revising the atomic model. Energy levels in atoms
What is the atomic model we have learned this far?
Who proposed this model?

3. Limitations of Rutherford's atomic model
Only explained a few simple properties of atoms.
Did not explain the chemical properties of atoms.
Ex. Couldn’t explain why metals or compounds of metals give off characteristic colors when heated in a flame.
Could not explain why an object such as the iron scroll first glows dull red, then yellow, and then white when heated to higher and higher temperature.
Being able to explain chemical properties of elements required a model that better described the behavior of the electrons in atoms.

4. The Bohr model In 1913 Niels Bohr developed a new atomic model
Experiment: Tested Hydrogen atom
Conclusion: Bohr proposed the an electron is found only in specific circular paths, or orbits around the nucleus.
Bohr proposed electrons orbit and have fixed energies which are called energy levels

5. Energy levels
Fixed energy levels of an electron are similar to rungs of a ladder
Electrons cannot exist between energy levels
To move from one energy level to the next an atom must gain or lose just the right amount of energy
A quantum of energy is the amount of energy required to move an electron from one energy level to the next

6. Energy levels
The amount of energy an electron gains or loses is not always the same
Higher energy levels are closer together
It takes less energy to move from one rung to the next near the top of the ladder
The higher the energy level occupied by an electron, the less energy it takes the electron to move from that energy level to the next higher energy level.
Failed to explain energies absorbed and emitted by atoms with more than one electron

7. Model:

8. The Quantum Mechanical Model
Bohr and Rutherford model of the atom describes the path of a moving electron
Experiment: Austrian physicist Erwin Schrodinger used these calculations and results to devise and solve a mathematical equation describing the behavior of the electron in a hydrogen atom.
Conclusion: theoretical calculations and experimental results were inconstant with describing the electron motion
Restricts the energy of electrons to certain values
Does not specify an exact path electron takes around the nucleus
Model: The quantum mechanical model came from the mathematical solutions to the Schrodinger equation which is the modern description of the electrons.

9. Cont…
In the quantum mechanical model, the probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloudlike region.
Probability describes how likely it is to find an electron in a particular location around the nucleus of an atom.
Cloud is dense and probability of finding an electron is high
For each energy level, the Schrödinger equation also leads to a mathematical expression, called an atomic orbital.

10. Review
How are the quantum mechanical model and the Bohr model alike? How are they different?

11. Like the Bohr model, the quantum mechanical model restricts the energy of electrons to certain values. Unlike the Bohr model, the quantum mechanical model does not specify an exact path the electron takes around the nucleus.

12. Atomic Orbitals Atomic Orbitals –
Describe the probability of finding an electron at various locations around the nucleus.
s orbitals – groups 1 and 2 in the periodic table.
p orbitals – groups 13-18
d orbitals – groups 3-12
f orbitals – lanthanide and actinide series

13. How do sublevels of principal energy levels differ?
Each energy sublevel corresponds to one or more orbitals of different shapes. The orbitals describe where an electron is likely to be found.

14. Different atomic orbitals are denoted by letters.
The s orbitals are spherical.
The p orbitals are dumbbell- shaped.
For a given principal energy level greater than 1, there is one s orbital, 3 p orbitals, and 5 d orbitals.
Four of the five d orbitals have the same shape but different orientations in space.

16. The numbers and types of atomic orbitals depend on the principal energy level.
Summary of Principal Energy Levels and Sublevels
Principal energy level
Number of sublevels
Type of sublevel
Maximum number 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

17. Orbitals
The principal quantum number, n, always equals the number of sublevels within that principal energy level.
The number of orbitals in a principal energy level is equal to n2.
A maximum of two electrons can occupy an orbital.
Therefore, the maximum number of electrons that can occupy a principal energy level is given by the formula 2n2.

18. Practice
How many orbitals are in the following sublevels? 3p 2s 4p 3d

19. 5.2 electron arrangement in atoms electron configuration
Energy and stability play an important role in determining how electrons are configured in an atom.
The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron configurations.
Three rules—the aufbau principle, the Pauli exclusion principle, and Hund’s rule—tell you how to find the electron configurations of atoms.

20. 1. Aufbau Principle
According to the aufbau principle, electrons occupy the orbitals of lowest energy first. In the aufbau diagram, each box represents an atomic orbital.
Increasing energy 6s 5s 4s 3s 2s 1s 6p 5p 5d 4p 4d 4f 3p 3d 2p
The aufbau diagram shows the relative energy levels of the various atomic orbitals. Orbitals of greater energy are higher on the diagram.
The range of energy levels within a principal energy level can overlap the energy levels of another principal level.

21. 2. Pauli Exclusion Principle
According to the Pauli exclusion principle, an atomic orbital may describe at most two electrons.
To occupy the same orbital, two electrons must have opposite spins; that is, the electron spins must be paired.
Spin is a quantum mechanical property of electrons and may be thought of as clockwise or counterclockwise.
A vertical arrow indicates an electron and its direction of spin ( or ).
An orbital containing paired electrons is written as

22. 3. Hund’s Rule
According to Hund’s rule, electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible.
Three electrons would occupy three orbitals of equal energy as follows.
Electrons then occupy each orbital so that their spins are paired with the first electron in the orbital.

23. Electron Configurations of Selected Elements Electron configuration
2px 2py 2pz 3s
Electron configuration H
1s1 He
1s2 Li
1s22s1 C
1s22s22p2 N
1s22s22p3 O
1s22s22p4 F
1s22s22p5 Ne
1s22s22p6 Na
1s22s22p63s1
An oxygen atom contains eight electrons.
The 1s orbital has two electrons of opposite spin.
The 2s orbital also has two electrons of opposite spin.
Each of the three 2p orbitals has one electron. The remaining electron now pairs with an electron occupying one of the 2p orbitals.

24. s p d f 1 2 3 4 5 6 7 1A 2A 3B 4B 5B 6B 7B 8B 1B 2B 3A 4A 5A 6A 7A 8A
group # = # valence (outside) e- s 1 2 3 4 5 6 7 p
Row =
# shells d f

25. Electron Configuration
1s1
group #
# valence e-
possibilities are:
s: 1 or 2
p: 1-6
d: 1-10
f: 1-14
Total e- should equal
Atomic #
row #
shell #
possibilities are 1-7
7 rows
subshell
possibilities are
s, p, d, or f
4 subshells
What element has an electron configuration of 1s1?

26. Practice: Ask these questions every time you have to write an electron configuration
Lithium:
find the element on the periodic table
what is the period number?
how many shells?
what is the group number?
how many valence electrons?
what subshell(s) does Li have?
what is the electron configuration?
atomic # = 3 2 2 1 1 s
1s2 2s1

27. Practice: Ask these questions every time you have to write an electron configuration
Boron:
find the element on the periodic table
what is the row #?
how many shells?
what is the group #?
how many valence electrons?
what subshell(s) does B have?
what is the electron configuration?
atomic # = 5 2 2 3 3 p
1s2 2s2 2p1

28. Subshells d and f are “special”
group # = # valence e- 1 2 3 4 5 6 7 3d d
period # = # e- shells 4d 5d 6d f 4f 5f

29. Electron Configuration
1s1
group #
# valence e-
possibilities are:
s: 1 or 2
p: 1-6
d: 1-10
f: 1-14
Total e- should equal
Atomic #
row #
shell #
possibilities are 1-7
7 rows
subshell
possibilities are
s, p, d, or f
4 subshells
What element has an electron configuration of 1s1?

30. Electron configuration
A convenient shorthand method for showing the electron configuration of an atom involves writing the energy level and the symbol for every sublevel occupied by an electron.
You indicate the number of electrons occupying that sublevel with a superscript.
For hydrogen, with one electron in a 1s orbital, the electron configuration is written 1s1.
For oxygen, with two electrons in a 1s orbital, two electrons in a 2s orbital, and four electrons in 2p orbitals, the electron configuration is 1s22s22p4.
Note that the sum of the superscripts equals the number of electrons in the atom.

31. practice
Write the electron configuration for:
A. P
B. C
C. Br
D. K

32. Short hand
To write short hand electron configuration use the Nobel gases because they are completely full.
Example:
Mg:
Al:
Na:

33. 5.3 atomic emission spectra
Waves
Wavelength: represented by  (the Greek letter lambda), distance from crest to crest
Short wavelengths = higher energy
Amplitude: distance from level of crest to zero
wavelength
amplitude

34. Waves
The frequency, represented by  (the Greek letter nu), is the number of wave cycles to pass a given point per unit of time unit the Hz or 1/s.
Higher frequency = higher energy

35. Cont…..
The product of frequency and wavelength equals a constant (c), the speed of light.
C is the speed of light  108 m/s
The frequency () and wavelength () of light are inversely proportional to each other. As the wavelength increases, the frequency decreases.
c = ln

36. All electromagnetic waves travel in a vacuum at the speed of light
Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays.
All electromagnetic waves travel in a vacuum at the speed of light
Low energy
( = 700 nm)
High energy
( = 380 nm)
Frequency  (s-1)
3 x 106
3 x 1012
3 x 1022
102
10-8
10-14

38. Pink floyd
When atoms absorb energy, their electrons move to higher energy levels. These electrons lose energy by emitting light when they return to lower energy levels.
A prism separates light into the colors it contains. White light produces a rainbow of colors.

39. Explanation of Atomic Spectra
Ground state
When an electron has its lowest possible energy, the atom is in its ground state.
When an electron is excited to a higher energy state it must absorb energy. When the electron goes back to its stable ground state it must emit energy.
The energy is emitted as a photon of light.

40. light
The energy absorbed by an electron for it to move from its current energy level to a higher energy level is identical to the energy of the light emitted by the electron as it drops back to its original energy level.
The wavelengths of the spectral lines are characteristic of the element, and they make up the atomic emission spectrum or Bright Line Spectra of the element.
No two elements have the same emission spectrum.
Hydrogen spectrum
Helium spectrum

42. practice
Calculate the wavelength of the yellow light emitted by a sodium lamp if the frequency of the radiation is 5.09 × Hz.

43. practice
What is the frequency of a red laser that has a wavelength of 676 nm?

44. photoelectric effect.
In the photoelectric effect, electrons are ejected when light shines on a metal.
The photoelectric effect could not be explained by classical physics.
Classical physics correctly described light as a form of energy.
But, it assumed that under weak light of any wavelength, an electron in a metal should eventually collect enough energy to be ejected.
Einstein’s theory that light behaves as a stream of particles explains the photoelectric effect and many other observations.
Light behaves as waves in other situations; we must consider that light possesses both wavelike and particle-like properties.
These light quanta are called photons.

45. The Photoelectric Effect
No electrons are ejected because the frequency of the light is below the threshold frequency.
If the light is at or above the threshold frequency, electrons are ejected.
If the frequency is increased, the ejected electrons will travel faster.