Modern PhysicsQuantum mechanics, Periodic Trends, atomic

orbitals, electron configuration,By Zoe Poncher and Abby Zemach

Quantum Mechanical Model

of the AtomSummary: 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

λ

De Broglie’s Equation

Speed of Light

Plank’s Constant Mass of

Electron

Avogadro's Number

History of AtomAntoine 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 ModelMax 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,

DefinitionsAtomic 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.

DefinitionsValence 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 TrendsSize dec.Ionization Energy inc.Electron Negativity inc.

Size inc.Ionization Energy dec.Electron Negativity dec.

Smallest Element

Largest Element

*Cations are usually SMALLER*

Electron Negativity

Mnemonic SkillDiagonal 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 negativityΔEN > 1.7 = Ionic Bond –inherently polar

Atomic Orbitals!S orbital

P orbital

D orbital

F orbital Electrons may be anywhere in the orbital,However, never in the origin!

Wherever the valance electrons lie from electron configuration shows what orbital the last electron Is in,

Electron ConfigurationElectron Configuration is the representation of

the arrangement of electrons

Full extended versionAll 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 versionUse 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’ orbitaln=2 & l=1

-1 0 1

Step Two – sketch orbital diagram and add number of valence electrons -Be Carful to Follow Aufbau’s and Hund’s Principals

0

Ml= 0 (lands in 0’s spot) &

Ms = ½ (I chose to establish arrow’s as

upwards first)

SourcesPeriodic 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