1. Electrons Galore!

2. Electrons and Light
By 1900, scientists knew that light could be thought of as moving waves that have given frequencies, speeds, and wavelengths.
In empty space, light waves travel at × 108 m/s.
The wavelength is the distance between two consecutive peaks or troughs of a wave.
The distance of a wavelength is usually measured in meters.
The wavelength of light can vary from 105 m to less than 10−13 m.
This broad range of wavelengths makes up the electromagnetic spectrum
Our eyes are sensitive to only a small portion of the electromagnetic spectrum.
This sensitivity ranges from 700 nm, which is about the value of wavelengths of red light, to 400 nm, which is about the value of wavelengths of violet light.

3. Electromagnetic Spectrum

4. Photoelectric Effect
In 1905, Albert Einstein proposed that light also has some properties of particles.
His theory would explain a phenomenon known as the photoelectric effect.
This effect happens when light strikes a metal and electrons are released.
What confused scientists was the observation that for a given metal, no electrons were emitted if the light’s frequency was below a certain value, no matter how long the light was on.
Yet if light were just a wave, then any frequency eventually should supply enough energy to remove an electron from the metal.
Einstein proposed that light has the properties of both waves and particles.
According to Einstein, light can be described as a stream of particles, the energy of which is determined by the light’s frequency.
To remove an electron, a particle of light has to have at least a minimum energy and therefore a minimum frequency.

5. Light is a wave
When passed through a glass prism, sunlight produces the visible spectrum
all of the colors of light that the human eye can see. You can see
Remember that the visible spectrum is only a tiny portion of the electromagnetic spectrum.
The electromagnetic spectrum also includes X rays, ultraviolet and infrared light, microwaves, and radio waves.
Each of these electromagnetic waves is referred to as light, although we cannot see these wavelengths.

6. Frequency, wavelength and Energy
Red light has a low frequency and a long wavelength.
Violet light has a high frequency and a short wavelength.
The frequency and wavelength of a wave are inversely related.

7. I see the light!
When a high-voltage current is passed through a tube of hydrogen gas at low pressure, lavender-colored light is seen. When this light passes through a prism, you can see that the light is made of only a few colors.
This spectrum of a few colors is called a line-emission spectrum.
Experiments with other gaseous elements show that each element has a line-emission spectrum that is made of a different pattern of colors.
In 1913, Bohr showed that hydrogen’s line-emission spectrum could be explained by assuming that the hydrogen atom’s electron can be in any one of a number of distinct energy levels.
The electron can move from a low energy level to a high energy level by absorbing energy.

8. I still see the light!
Electrons at a higher energy level are unstable and can move to a lower energy level by releasing energy.
This energy is released as light that has a specific wavelength.
Each different move from a particular energy level to a lower energy level will release light of a different wavelength.
Bohr developed an equation to calculate all of the possible energies of the electron in a hydrogen atom.
His values agreed with those calculated from the wavelengths observed in hydrogen’s line-emission spectrum.
In fact, his values matched with the experimental values so well that his atomic model that is described earlier was quickly accepted.

9. Light tells me where electrons live!
Normally, if an electron is in a state of lowest possible energy, it is in a ground state
If an electron gains energy, it moves to an excited state
An electron in an excited state will release a specific quantity of energy as it quickly “falls” back to its ground state.
This energy is emitted as certain wavelengths of light, which give each element a unique line-emission spectrum

10. Electron Transitions

11. Quantum Numbers
The present-day model of the atom, in which electrons are located in orbitals, is also known as the quantum model.
According to this model, electrons within an energy level are located in orbitals, regions of high probability for finding a particular electron.
However, the model does not explain how the electrons move about the nucleus to create these regions.
To define the region in which electrons can be found, scientists have assigned four to each electron.

12. Orbitals

13. n tells you in which period the atom is
The principal quantum number, symbolized by n, indicates the main energy level occupied by the electron.
Values of n are positive integers, such as 1, 2, 3, and 4.
As n increases, the electron’s distance from the nucleus and the electron’s energy increases.

14. l tells you the shape of the orbital
The main energy levels can be divided into sublevels.
These sublevels are represented by the angular momentum quantum number, l.
This quantum number indicates the shape or type of orbital that corresponds to a particular sublevel. There are letter codes for this quantum number
l = 0 corresponds to an s orbital
l = 1 to a p orbital
l = 2 to a d orbital
l = 3 to an f orbital.
For example, an orbital with n = 3 and l = 1 is called a 3p orbital, and an electron occupying that orbital is called a 3p electron.

15. s Orbitals (l = 0)—Shape Each shell has one s orbital.
Lowest energy orbital in a shell.
Spherical in shape.

16. s Orbitals (l = 0)—Shape For n = 2 and l = 0 For n = 3 and l = 0
The larger the value of n, the larger the orbital.
The further out from the nucleus it extends.
Contains nodes-regions of no probability.
number of nodes = (n – 1)

17. p Orbitals (l = 1)—Shape
Each principal energy state n = 2 and above has three p orbitals.
ml = -1, 0, +1
Each of the 3 orbitals point along a different axis.
px, py, pz
2nd lowest energy orbitals in a principal energy state.
Shape has two lobes.
Node at the nucleus. Total of n nodes.

18. d Orbitals (l = 2)—Shape
Each principal energy state n = 3 and above has five d orbitals
ml = -2, -1, 0, +1, +2
4 of the 5 orbitals are aligned in a different plane
the fifth is aligned with the z axis, dz squared
dxy, dyz, dxz, dx squared – y squared
3rd lowest energy orbitals in a principal energy state
Mainly 4-lobed.
one is two-lobed with a toroid

19. f Orbitals (l = 3)—Shape
Each principal energy state n = 4 and above has seven f orbitals
ml = -3, -2, -1, 0, +1, +2, +3
4th lowest energy orbitals in a principal energy state.
Mainly 8-lobed
some 2-lobed with a toroid

20. mL tells you the orientation
The magnetic quantum number, symbolized by m, is a subset of the l quantum number.
It also indicates the numbers and orientations of orbitals around the nucleus.
The value of m takes whole-number values, depending on the value of l.
The number of orbitals includes one s orbital, three p orbitals, five d orbitals, and seven f orbitals.

21. ms tells you the magnetic nature
The spin quantum number is symbolized by ±½
(↑ or ↓), indicates the orientation of an electron’s magnetic field relative to an outside magnetic field.
A single orbital can hold a maximum of two electrons, which must have opposite spins.

22. Electron Configurations
Each orbital can hold a maximum of two electrons.
The discovery that two, but no more than two, electrons can occupy a single orbital was made in 1925 by the German chemist Wolfgang Pauli.
This rule is known as the Pauli Exclusion Principle
Another way of stating the Pauli exclusion principle is that no two electrons in the same atom can have the same four quantum numbers.
The two electrons can have the same value of n, the same value of l and they may
have the same value of m. But these two electrons cannot have the same spin quantum number.
If one electron has the value of +½ then the other electron must have the value of -½

23. Electron Configurations
The arrangement of electrons in an atom is usually shown by writing an electron configuration.
Like all systems in nature, electrons in atoms tend to assume arrangements that have the lowest possible energies.
An electron configuration of an atom shows the lowest-energy arrangement of the electrons for the element.

24. Electron Configuration of Atoms
Every electron in an atom has a unique set of quantum numbers.
Quantum numbers describe which orbitals are occupied by electrons in an atom.
Only two electrons can occupy one orbital.
Arrangement of electrons an atom is called the electron configuration.
lowest energy arrangement = ground state

25. Electron Configuration of Atoms
Two ways to depict the electronic structure of an atom.
Orbital filling diagram.
Follow Aufbau Principle.
Electron configuration.
Full electron configuration.
Condensed using the “inner electron configuration” based on the closest previous Noble Gas.

26. Rules of the Aufbau Principle
Determine number of electrons in atom or ion.
Orbitals are filled in order of increasing energy, lower energy orbitals first.
Lowest energy shells fill first.
n =1 before n = 2, etc.
Lowest energy orbitals in a shell fill first.
s  p  d  f
An orbital can hold two electrons of opposite spin—Pauli Exclusion Principle.
For degenerate orbitals (same energy), the lowest energy is attained when the number of electrons with the same spin is maximized—Hund’s Rule.
Half fill orbitals of the same energy before pairing electrons.
One electron per orbital until subshell is half full.
All have same spin, prior to pairing.

27. Electron Configuration Example1s 2s 2p 3s 3p 3d 4s 4p 4d 5s 5p Na
Atomic Number (Z): 11
1s2 2s2 2p6 3s1 or [Ne] 3s1
n = 1 l = 0
n = 2 l = 0
n = 2 l = 1
n = 3 l = 0
n = 3 l = 1

28. Electron Configuration More
Identify the element with the ground state configuration
1s2 2s2 2p4
Carbon
Oxygen
Sulfur
Argon

29. Electron Configuration Your Turn
Draw the ground state orbital filling diagram for vanadium.
[Ar]  _ _ _
[Ar] _ _ _ _ _ __
[Ar] _ _ _ _ _
[Ar]  _ _ _ __ __ 4s 4p 4s 4p 3d 3d 4s 3d

30. Electron Configuration and the Periodic Table
s block
p block
d block
f block
Begin
here
End

31. Aufbau Principle
Begin
here
End
here

32. Electron Configuration and the Periodic Table
The group number corresponds to the number of valence electrons.
The Period number corresponds to the principle energy level (n) of the valence electrons.
The length of each “block” is the maximum number of electrons the subshell can hold.
Subshell = same orbital type in a level
Example: all p orbitals (l = 1) in n = 3 level.

33. More Electron Configuration
What is the ground state electron configuration for germanium (Ge)?
[Ar] 3d10 4s2 4p2
[Ar] 4s2 4p6 3d6
[Ar] 4s2 3d10 4p2
[Ar] 3d10 4s2 4p1

34. More Electron Configuration
What is the ground state electron configuration for tantalum (Ta)?
[Xe] 6s2 5d3
[Xe] 6s2 4f14 5d3
[Xe] 6s2 5d10 4p7
[Xe] 5d10 6s2 4p7