1. Electrons in Atoms Chapter 5
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2. Section 5.1 Objectives
Define a quantum of energy, and explain how it is related to an energy change of matter.
Compare the wave and particle natures of light
Contrast continuous electromagnetic spectra and atomic emission spectra.
Review
Radiation: the rays and particles —alpha particles, beta particles, and gamma rays—that are emitted by radioactive material
The Rutherford’s model doesn’t explain why negatively charged electrons aren’t pulled into the positively charged nucleus

3. Electromagnetic Radiation
Visible light is a type of electromagnetic radiation, a form of energy that exhibits wave-like behavior as it travels through space
Radio waves, microwaves, visible light, and x rays are all examples of electromagnetic waves that differ from each other in wavelength and frequency
Most subatomic particles behave as Particles and obey the physics of waves

4. Wave Characteristics Hz = 1/sec
Frequency () – number of waves that pass a point per second
Units are “cycles per sec” or Hertz
Hz = 1/sec
Wavelength () – distance from crest to crest
Amplitude – is the wave’s height from the origin to a crest
All radiation: c =  • 
c = velocity of light = 3.00 x 108 m/sec
Crest
Amplitude
Wave Cycle
Trough

5. Electromagnetic Spectrum
White light (sunlight) is made up of Continuous Spectrum (7 colors of rainbow)
ROY G BIV = Red, Orange, Yellow, Green, Blue, indigo, violet
Each color has a specific Frequency & wavelength
Electromagnetic spectrum

6. Wavelength vs. Frequency
As frequency increases the wavelength decreases (indirectly proportional)
Long wavelength --> small frequency --> low energy
Short wavelength --> high frequency --> High energy

7. Bright-line Spectrum
Bright-line Spectrum – distinct separate lines of color for an element
Atoms emit an atomic emission spectra, characteristic set of discrete wavelengths
not a continuous spectrum
Atomic Emission spectrum can be used as a “fingerprint” for an element

8. H, Hg, and Ne Spectrums
Light can be grouped into packets called Photons
Like matter is divided into atoms
Energy of Photon uses the eqn: E = h 
E = energy (Joules, J)
H = Planck constant x J·sec

9. Section 5.2 & 5.3 Objectives
Identify the relationship among an atom’s energy levels, sublevels, and atomic orbital’s
Apply the Pauli exclusion principle, the aufbau principle, and Hund’s rule to write electron configurations using orbital diagrams and electron configuration notation
Define Valence electrons, and draw electron-dot structures representing an atom’s valence electrons

10. Terms
The wave model of light cannot explain all of light’s characteristics
A beam of light has wave like and particle like particles
Matter can gain or lose energy only in small, specific amounts called quanta
A quantum is the minimum amount of energy that can be gained or lost by an atom
A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy

11. Electron Configuration
Is the arrangement of electrons around the nuclei
A listing showing how many electrons are in each orbital or subshell in an atom or ion
Example:
Bohr atomic model/Quantum Mechanics
Bohr's model was rather simplistic and as scientists made more discoveries about more complex atoms, Bohr's model was modified and eventually was replaced by more sophisticated models.
The Quantum Mechanical Model of the atom presents a more accurate model of the atom. It is a more sophisticated model based on complex mathematical calculations and interpretations
1s s p s1

12. Electron Configuration
Quantum mechanics is the study of the relationship between energy quanta (radiation) and matter, in particular that between valence shell electrons and photons
Electrons in atoms are arranged (or located) as:
LEVELS (n)
Energy level of electron – region electron is likely to be found
SUBLEVELS
ORBITALS

13. Bohr’s Model of the Atom
Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits
Energy Added
(A quantum)
Hydrogen’s single electron is in the n=1 orbit in the ground state
When energy is added, the electron moves to the n=2 orbit

14. How are photons emitted?
Electron starts at 2nd energy level
grounded state
Excited electron steps from current to higher level
Quantum energy is the amount of energy required to move an electron from its present energy level (grounded state) to the next higher level (excited state)
Energy is absorbed and released by atoms in certain fixed amounts known as quanta
quanta = bundles of energy
Excited State
n=3
Quantum Energy
Grounded state
n=2
n=1

15. How are photons emitted?
Electron at excited state
Electron will “fall” back to 2nd energy level
Photon released
At a specific color
Due to frequency and wavelength
The photon is equal in energy to the quantum energy added
Excited State
n=3
Grounded state
n=2
n=1

16. How are photons emitted?
Electron starts at 2nd energy level
grounded state
3. Excited again
This time adding more Quantum energy
Result: higher energy level reached
4. When electron “falls”
Photon different color
Higher frequency, shorter wavelength
More energy = higher energy color
Excited State
n=4
Quantum Energy
n=3
Grounded state
n=2
n=1

17. How are photons emitted?
We can only see electrons that start at a grounded state at level 2.
Why??
We can only see a certain wavelength and frequency
Electrons returning to other levels do emit frequency’s but they are out of the visible light spectrum
n=3
n=2
n=1

18. Electron Configurations
Sublevel
2p4
Number of electrons in the sublevel
Energy Level
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14… etc.

19. Principal Quantum Numbers
1s s p s1
Energy Levels are called Principal quantum numbers (n)
Assigned values in increasing energy:
n=1,2,3,4,5,6,7
(think of n as an address for electrons)
corresponds with rows
Energy Level n
Max # of electrons 1 2 8 3 18 4 32 5
n = 1
n = 2
n = 3
n = 4

20. Sublevels 1s2 2s2 2p6 3s1 Sublevels are in each Energy level
# of sublevels equals the principal quantum number (n)
Max # of sublevels is 4
The sublevels are:
s, p, d, f
Sublevel
Max # of electrons s 2 p 6 d 10 f 14

21. Orbitals 1s2 2s2 2p6 3s1 In each sublevels there are orbitals:
S subshell (1 orbital)
P subshell (3 orbitals)
d subshell (5 orbitals)
F subshell (7 orbitals)
Each orbital holds 2 electrons, that spin in opposite direction
Total # of orbitals in a n level n2
Total # of electrons in a n level
2n2
An atomic orbital is a mathematical function that describes the wave-like behavior of an electron in an atom.

22. Orbitals and the Periodic Table
Orbitals grouped in s, p, d, and f orbitals
s orbitals
d orbitals
p orbitals
f orbitals

23. Rules for electron configuration
1s s p s1
Aufbau Principle
Electrons enter orbitals of lowest energy first, starting with the 1s orbital
Pauli Exclusion Principle
An orbital is filled when it contains 2 electrons.  After that, you have to put the electrons in a different orbital
Hund’s Rule
Within a subshell, the electrons will occupy the orbitals singly first, and will only pair up when there are no longer any empty orbitals available in that subshell

24. Valence Electrons
Valence Electrons are defined as electrons in the atom’s outermost orbitals – those associated with the atoms highest principal energy level
Electron-dot structure consists of the elements symbol representing the nucleus, surrounded by dots representing the element’s valence electrons
The octet rule states that atoms with 8 electrons in their outer shell are stable.

25. Extra Info

26. An excited lithium atom emitting a photon of red light to drop to a lower energy state.

28. Orbital occupancy for the first 10 elements, H through Ne.

29. Order to fill orbital’s

30. Aufbau principle
The Aufbau principle states that each electron occupies the lowest energy orbital available

31. Dot Diagram order for Valence electrons
Below you will see an example of the order of filling in the dots on a dot diagram for an element with eight valence electrons.   Please note that you can place the first two dots on any side, but the rest of the dots should be placed in either a clockwise or counter clockwise manner, with no side receiving two dots until each side gets one.
Step 1
Step 2
Step 3
Step 4
Step 5
Step 6
Step 7
Step 8

32. Quantum Mechanical Model
Quantum numbers are used to describe certain aspects of the locations of electrons.
The shape, size, and energy of each orbital is a function of 3 quantum numbers
n (principal quantum number)
energy level, position of the electron with respect to the nucleus
Sublevels
shape of orbital
s,p,d,f
Orbitals
describes the direction of the electron’s spin within a given orbital (clockwise or counterclockwise)
1s s p s1