1. Light, Quantitized Energy & Quantum Theory EQ: What does the Modern Atom look like?
CVHS Chemistry Ch 5

2. Problems w/ Rutherford’s model
Didn’t explain:
Where Electrons were
Why Electrons didn’t fall into + charged Nucleus
Chemical Behavior of different elements

3. Flame Tests

4. Wave Nature of Light
Amplitude: wave height from origin (vertical center) to crest or trough
Wavelength (λ)
Measured from crest to crest or trough to trough
The wavelength of light is usually expressed in nanometers (1 nm = 1 x 10–9 m).

5. Wave Nature of Light Frequency (n): nu
# of waves that pass a given point per second Hz (/s)
In calculations, frequency is expressed with units of “waves per second,”
or (s–1) where the term “waves” is understood

6. Plank & The Quantum Concept
matter can gain or lose energy only in small, specific amounts called quanta.
That is, a quantum is the minimum amount of energy that can be gained or lost by an atom.

7. Photoelectric Effect
In the photoelectric effect, electrons, called photoelectrons, are emitted from a metal’s surface when light of a certain frequency shines on the surface.
Further, Einstein proposed The photon of light must have enough energy to bump the electron off the metal

8. Atomic Emission Spectra
The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by atoms of the element
Each emission spectra is unique, like a fingerprint

9. Bohr’s Model of the Atom
Since Energy is quantitized (Planck & Einstein) atoms only have 1 allowable energy state
Ground State: Lowest Energy state
Excited State: Energy added above ground state
Energy state is related to location of electron around nucleus
Ground state: Close to nucleus
Excited state: Away from nucleus
Farther away = More Energy

10. Louis de Broglie’s Atom
Treated electrons orbit as a wave
NOTICE: Waves have mass!

11. Heisenberg Uncertainty Principle
Measuring the location or velocity of an electron changes the location and velocity of the electron
The Heisenberg uncertainty principle Can’t know velocity & same time

12. Physicist Erwin Schrödinger (1926) & Wave Equations
Treated the electron as a wave
Aka: quantum mechanical model
Seemed to apply to all atoms, not just hydrogen like Bohr’s model
Puts electrons in distinct energy levels but doesn’t describe its path around the nucleus
Atomic Orbital: 3-D region that describes an electrons most probable location around the nucleus
Sort of a fuzzy cloud
The circle represents the 90% probability area for finding an electron

13. Quantum Numbers Principal Quantum Number (n)
Distance from the nucleus
The # of sublevels found in an energy level is equal to the number of the energy level
Energy Sublevels: describe by the shape of the atom’s orbitals
s, spherical (1)
p, peanut or dumbell (3)
d, daisy (variable) (5)
f, flower (variable) (7)

14. Energy Levels & Sublevels
Electrons w/in the same energy level but in different sublevels have similar amounts of energy
Sublevels
s, p, d, f
s has the lowest energy & f has the highest energy
Each energy level has the # of sublevels equal to the number of the energy level
Energy level 1 has 1s
Energy level 2 has
2s & 2p
Energy level 3 has
3s & 3p & 3d
Which orbital would it take the most energy to travel around?
F because it’s so complex

15. Electron Configurations
A specific # of electrons can fit into each sublevel s
Holds 2 electrons p
Holds 6 electrons d
Holds 10 electrons f
Holds 14 electrons

16. Rules for Electrons
Aufbau Principle: Electrons fill the lowest energy levels first
Hund’s Rule: At any sublevel, electrons go into orbitals one at a time until they have no choice but to pair.
Pauli Exclusion Principle: Any 2 e- in the same orbital must spin in opposite directions.

17. Electron Configurations
Electrons fill energy levels and sublevels in an orderly fashion w/ the lowest energy orbitals being filled first
This creates the most stable electron configuration
The shape of the modern periodic table is a direct result of the arrangement of electrons in the atom.

18. Electron Configurations
The periodic table is divided into blocks that show the sublevels and orbitals occupied by the electrons of the atoms.
n represents the energy level and the sublevels are shown as well.
Notice how d and f sublevels have lower energy levels than other atoms in their rows, this is because the pattern that the electrons have to move in is more complex than the s and p sublevels. So an atom will put electrons in a higher energy level s or p sublevel before it puts any electrons in the d sublevel

19. Blank Periodic table
Take a moment to fill to fill out a blank periodic table

20. How to write an electron configuration
1. Find the symbol for the element
2. Write the symbol in brackets for the nearest smaller noble gas
3 Write the outer electron configuration for the remaining electrons.

21. Practice Electron ConfigurationsLi Be B C N O F Ne

22. Periodicity
Write Noble gas configurations for: Be Mg Ca Sr Ba Ra
Elements in same group or family have similar chemical properties
Elements in same period have similar configurations
Periodic table = periodicity: predictable patterns based on electron configurations.

23. Transition Elements
Notice in the periodic table that calcium is followed by a group of ten elements beginning with scandium and ending with zinc
These elements begin to fill the 3d sublevel
d is very complex so electrons fill 4s and then 3d
Metals so they lose electrons and become cations

24. Inner Transition Elements
Two rows at bottom of periodic table
lanthanides (atomic numbers 58 to 71)
actinides (atomic numbers 90 to 103)
These two series are called inner transition elements because their last electron occupies inner-level 4f orbitals in the sixth period and the 5f orbitals in the seventh period

25. An easy way to remember how atoms fill orbitals & sublevels