1. By Zoe Poncher and Abby Zemach
Modern Physics
Quantum mechanics, Periodic Trends, atomic orbitals, electron configuration,
By Zoe Poncher and Abby Zemach

2. Quantum Mechanical Model of the Atom
Summary: Theory that electrons have both wave like and particle properties, and reside in orbitals (each of which hold two electrons).
Names:
Heisenberg – Uncertainty Principle – both position and momentum cannot be known at the same time
Louis De Broglie – De Broglie Eq = λ = h/mv
Erwin Schrodinger—created equation that showed the quantum state, or wave function, of a particle as its position or time interval changed
Max Planck- E = hv=h(c/λ): Planck’s Constant = 6.63χ10-34

3. Equations De Broglie’s Equation λ Avogadro's Number Speed of Light
Plank’s Constant
Avogadro's Number
Mass of Electron

4. History of Atom Antoine Lavoisier – 1774 Law of Conservation of Matter
Joseph Proust – Law of Constant Composition
John Dalton - Atomic Theory & Law of Multiple Proportions
Michael Faraday-demonstrated electric nature of elements
Sir William Crookes – Cathode ray tube
JJ Thomson – Discovered Electrons- oil drop experiment
Ernest Rutherford – nuclear model of atom w/ Gold Foil exp.
Discovered Proton

5. Quantum Model E = hv: Planks Constant:6.62606957 × 10-34 J/ s
Max Planck – German physicist ( ) hypothesized that energy could be released or absorbed by atoms in discrete “chunks” of minimum size . ‘Quantum meant the smallest quantity of energy that could be emitted or absorbed as electromagnetic radiation.
E = hv:
Planks Constant: × J/ s

6. Photoelectric Effect Ephoton –Ethreshold = KEelectron
Albert Einstein used the Quantum theory to explain why energy acts like a particle when it impacts metal, compared to wave property- otherwise known as the Photoelectric Effect!
Ephoton –Ethreshold = KEelectron
Ephoton= hf = hc/λ

7. Principles You NEED to Know
Pauli Exclusion Principle –in a given atom no two electrons can have the same quantum #’s / can’t have same spin direction.
Aufbau Principle-electrons fill lowest energy orbitals first,
Hund’s Rule- the lowest energy state- which is the most stable –is the one with the greatest # of unpaired electrons,

8. Definitions
Atomic size: Atomic radii between nuclei and outer electron shell.
Ionization Energy: Energy required to remove an electron from an atom in the gas phase
Electron Negativity: The ability of an atom to attract electrons,
*** Trends develop from the distance which the outer shell is from the nuclei, and the ratio of protons to electrons. Ie Oxygen is smaller than Nitrogen because the effective nuclear charge is greater (more protons) increasing the pull on the outer orbitals.

9. Definitions
Valence Electrons – Outermost electrons, Isoelectric – Atoms having same electron configuration Paramagnetic- Atoms with UNPAIRED electrons which are attracted to magnetic fields, Diamagnetism – contains only paired electrons ,

10. Periodic Trends *Cations are usually SMALLER* Size dec.
Ionization Energy inc.
Electron Negativity inc.
Smallest Element
Largest Element
Size inc.
Ionization Energy dec.
Electron Negativity dec.
*Cations are usually SMALLER*

11. Electron Negativity Mnemonic Skill Diagonal Trend:
Estimate Electron Negativities by knowing EN increase as you go up and across the periodic table.
Electron Negativity Differences and Polarities
ΔEN <.4 = Covalent Bond – non polar
.4 < ΔEN <1.7 = < Polar Covalent – Polar w/ Dipole pointing to higher negativity
ΔEN > 1.7 = Ionic Bond –inherently polar

12. Atomic Orbitals! S orbital P orbital D orbital F orbital
Wherever the valance electrons lie from electron configuration shows what orbital the last electron
Is in,
S orbital
P orbital
D orbital
F orbital
Electrons may be anywhere in the orbital,
However, never in the origin!

13. Electron Configuration
Electron Configuration is the representation of the arrangement of electrons
Full extended version
All of the periods, orbitals, and electrons that are to the specific element, ie Se
Se: 1s2 2s2 2p6 3s2 3s6 4s2 3d10 4p4
The principal energy levels
The orbitals used
Number of electrons
filled in the orbital
Abbreviated version
Use previous Noble gas as place holder for all periods and orbitals in previous periods
Se: [Ar] 4s2 3d10 4p4

14. Electron Configuration
Please Note Electrons which are Excited jump to a higher energy level, meaning electrons are promoted,
When you promote electrons you take from the inner orbitals first,
Before
After
When you remove electrons you take from the outermost shell,            

15. Quantum Numbers s =0, p=1 , d=2 , f=3
n = Principal Quantum # = Period / energy level
l = Angular Momentum (azimuthal) Quantum # =
s =0, p=1 , d=2 , f=3
Ml = Magnetic Quantum # = - l - l value,
ie. P orbitals have -1,0,1 #’s
Ms = Spin Quantum # = -½ , ½ depends on electron configuration
** To find the Quantum numbers you use a combination of looking at the Periodic table and orbital diagram**
No two different elements have the same set of Quantum Numbers

16. Quantum Numbers
What are the Quantum Numbers for Carbon?
Step One – notice Carbon is in the
2nd Period and P’ orbital
n=2 & l=1  -1 1 
Step Two – sketch orbital diagram
and add number of valence electrons
-Be Carful to Follow Aufbau’s and Hund’s
Principals
Ml= 0
(lands in 0’s spot)
&
Ms = ½ (I chose to establish arrow’s as upwards first)

17. Sources Periodic Table Diagram: http://www.chemicalelements.com/
AP Chemistry Crash Course: by Michael D’Alessia,
Barron’s AP Chemistry Exam Book
Chemistry The Central Science 11e