1. Energy & Electrons

2. Energy
One way energy travels through space is by electromagnetic radiation.
This energy has wave-like properties
It has wavelength (λ) UNITS: m
It has frequency (υ) (the number of cycles passing in a given time (usually seconds) UNITS: /s OR Hz
It has amplitude

4. Energy It also has speed.
In a vacuum, all electromagnetic radiation (EMR)travels at the speed of light (c )
When passing through matter, its speed decreases.
Due to the fact that there is very little matter in air, EMR travels through air also at the speed of light (c )
c = x 108 m/s

5. Waves Since all waves in the Electromagnetic Spectrum
have the same speed,
c = λ υ
This means that wavelength and frequency are inversely proportional:
a long wavelength has a low frequency
a short wavelength has a high frequency
1. In the Electromagnetic Spectrum, what has the longest wavelength?
2. In the Electromagnetic Spectrum, what has the shortest wavelength?
3. In the Electromagnetic Spectrum, what has the highest frequency?

6. Dual Nature of Light
Light exhibits both wave properties and characteristics of particulate matter.
Light is “quantized” (it is made up of photons)
Photons (the stream of particles in a wave) have mass-like characteristics when in motion.

7. Max Planck
Max Planck found that matter could NOT absorb and emit any quantity of energy.
Instead, energy is lost and gained in whole numbers of hυ
Each 1 of these is (hυ) is called a “quantum”
Black body radiation

8. The Photoelectric Effect
Einstein noted that when Electromagnetic Radiation strikes the surface of metals, electrons were given off.
In order for this to occur the incident light hitting the surface has to reach a threshold frequency.
The lowest frequency required (the threshold frequency) was calculated
E = h υ
E is energy of the photon
h is Planck’s constant (6.626 x J

9. Atomic Emission Spectrum
When white light passes through a prism, it is broken into its parts
When electricity passes through a gas, it causes lines of color to appear.
This is called the “emission spectrum”

10. The Emission (Line) Spectrum
Each element has a specific emission spectrum
It can be used to help identify the elements present in an unknown
The line spectrum also informs us that only certain energies are allowed for certain electrons

11. The Bohr Model
Niels Bohr proposed that electrons are held in circular orbits.
When an electron is stuck by energy with a high enough frequency, it jumps from its normal (ground state) position, to an excited state.
When the electron goes back to its state of rest, it gives off energy in the form of visible light.

12. Hydrogen’s Emission Spectrum

13. Emission Spectra
The color depends on how the energy level the electron has moved to.

14. Calculating Energy
En = -RH (z2) (n2) RH = Rydeberg constant = 2.18 x 10-18J Z = nuclear charge n = energy level

15. Calculating Energy
In Hydrogen, when an electron is in the 6th energy level
n = 6
E6 = -(2.18 x 10-18J) (12)
(62)
E6 = x 10-20J
the photon has lost x 10-20J to the electron (the electron absorbs it)

16. Calculating Energy Emitted
When an electron of hydrogen drops from the 6th energy level to the 1st energy level, a photon is given off. Calculate its energy.
(remember: exponents have to match to add or subtract)
ΔE = En(final) – En(initial)
ΔE = E1 – E6

17. Calculating Energy required
Due to the Conservation of Energy,
The energy emitted is the same as energy required
These equations are specific to hydrogen because the Bohr model does not accurately describe electron position.
However, the emission spectra are still useful.

18. Based on the equation E = h υ and c = λ υ We can solve for E in terms of h, c, and λ E = h c λ We can use this to calculate the wavelength of a photon emitted when an electron returns to its ground state.

19. More electrons in an atom will result in more lines

20. Determining Electron Position in the Atom
All electrons have a specific energy level, sublevel, alignment, and spin.
Energy level (n) the relative distance from the nucleus.

21. Sublevels Sublevel (l) The shape of the path an electron travels in.
The sublevels are s, p, d, and f
An s sublevel can hold up to 2 electrons
A p sublevel can hold up to 6 electrons
A d sublevel can hold up to 10 electrons
An f sublevel can hold up to 14 electrons

22. Orbitals
Describes the relative distance from the nucleus, the path, and its alignment
Each orbital can hold 2 electrons
There are 3 p orbitals (6 electrons total/2 electrons each orbital)
For example 2px is an orbital
Due to the fact that orbitals of the same energy level and sublevel have the same energy, they are often lumped together
For example 2px 2py and 2pz
make up the 2p orbitals

23. S and P orbitals
Valence electrons are located in the s and p orbitals

24. d orbitals

25. f orbitals

26. Aufbau Principle: Electrons fill the orbitals with the lowest energy first.

27. Using the above Example Cl p = 17
e = 17 (because this Cl has no charge
Start at 1s (it can hold up to 2, 17-2=15 left)
1s2
Then continue to 2s (it can hold up to 2, 15-2=13 left)
1s22s2
Then continue to 2p (it can hold up to 6, 13-6 = 7 left)
1s22s22p6
Then continue to 3s (it can hold up to 2, 7-2 = 5 left)
1s22s22p63s2
Then continue to 3p (it can hold up to 6, only 5 left
1s22s22p63s23p5
(the last step leaves you with the electron configuration for Cl)

29. To use the above
For Cl Locate Cl on the periodic table Move left to right, top to bottom on the periodic table until you reach Cl 1s2 (you pass through 2 boxes) 2s2 (you pass through 2 boxes) 2p6 (you pass through 6 boxes) 3s2 (you pass through 2 boxes) 3p5 (only 5 boxes to get to Cl) The final electron configuration is, 1s22s22p63s23p5