Modern Physics
Quantum mechanics, Periodic Trends, atomic orbitals,
electron configuration,
By Zoe Poncher and Abby Zemach
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
Equations
De Broglie’s
Equation
λ
Speed of
Light
Plank’s
Constant
Mass of Electron
Avogadro's
Number
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
Quantum Model
Max Planck – German physicist (1858-1947)
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:6.62606957 ×
10-34 J/ s
Photoelectric Effect
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
Principles You NEED to
Know
Principles:
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,
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.
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 ,
Periodic Trends
Size dec.
Ionization Energy inc.
Electron Negativity inc.
Smallest Element
Largest Element
Size inc.
Ionization Energy dec.
Electron Negativity dec.
*Cations are usually SMALLER*
Electron Negativity
Electron Negativity Differences and Polarities
Mnemonic Skill
Diagonal Trend:
Estimate Electron
Negativities by knowing EN
increase as you go up and
across the periodic table.
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
Atomic Orbitals!
S orbital
Wherever the valance
electrons lie from electron
configuration shows what
orbital the last electron
Is in,
P orbital
D orbital
F orbital
Electrons may be anywhere in the orbital,
However, never in the origin!
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
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,
Before





After


Quantum Numbers
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
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
0
-1
0



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)
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