Chemistry I Honors--Unit 3: Quantum Mechanical Model of the

Atom & Periodic Trends

Objectives #1-7: The Development of a New Atomic

Model I. Electromagnetic RadiationIn the late 1800’s and early 1900’s,

scientists discovered that passing an electric current through gases of various elements caused electromagnetic radiation in the form of colored light to be emitted from the gas.

Examples of colors produced by the electrically charged gasses include:

Sodium YellowCopper GreenNeon Orange

Hydrogen Purple

Additional testing showed that EMR of energies too low or too high to see with the eye were also produced

Electromagnetic radiation is energy that travels in the form of a wave

Examples of Electromagnetic Waves

All waves have AMPLITUDE ,FREQUENCY & WAVELENGTH

Characteristics of EMR:The wavelength of a wave is the distance between the peaks of the waveThe frequency of a wave is the number of peaks that pass by a point is space in one second (the rate of reproducibility)The speed of all EMR is the speed of light (3.00 X 108 m/s)

Wavelength and frequency are inversely related to each other

c = (f) ()Frequency and energy are directly related to

each otherE = (h) ()

Wavelength and energy are inversely related to each other

E = (h)(c) / ()h= 6.626 x 10-34 J . Sec

These relationships were discovered by Max Planck.

Types of EMR

Lower Energy Higher Energy

Radio Radar Micro IR Visible* UV X-rays Gamma

red, orange, yellow, green, blue, indigo,

violet

Objectives #1-7: The Development of a New Atomic Model

I. Electromagnetic Radiation* EMR refers to all the various types of

radiant energy, from radio waves to gamma waves

II. The Origins of Wave Mechanics

EMR has dual qualities: (Louis DeBroglie, 1892-1987), French

“EMR is like 4 year old who can’t make up

their mind what to be for

Halloween!! “A wave! No,

wait, a particle! No wait, a wave…

no wait…..”

EMR acts as a particle when it interacts with matter; this is illustrated by the photoelectric effect which involves the emission of electrons when radiation of a specific frequency strikes the surface of a metal (Albert Einstein, 1905, German-American)…

= (h) (v0)EMR acts as a wave when it travels

through space… = (h) /(m) (v)

Photoelectric Effect

Incoming energy waves Electrons emitted

Video Clip: The Photoelectric Effect

Albert Einstein, 1905, German-American

“The photoelectric effect helps

explain why your solar calculator

works!”

Wave Particle DualityDe Broglie discovered that EMR acts as a wave when it travels through space… = (h) /(m) (v)

The Double Slit Experiment

• Evidence of wave particle duality of electrons…

• Video Clip: “The Infamous Double Slit

Experiment”

Objectives #1-7 The Development of a New Atomic Model

Atomic Emission Spectrum: an explanation of the colors produced by exciting atoms

Niels Bohr, 1885-1962, Danish

“Electrons LOVE to jump energy levels —they emit light as

they move back ‘home’!!”

Examples of Atomic Emission Spectrums

Bohr’s Theory of Light Emission

An electron is normally in its low energy state or ground state.

When the electron becomes excited with a certain amount of energy or quantum, it will “jump” to a higher level of energy or its excited state.

This new state is unstable for the electron and so this excess energy is emitted as a photon of EMR and the electron returns to the ground state.

III.The Bohr Model of the Atom Proposed that electrons revolve around

the nucleus in definite paths or orbits Each electron has a certain amount of

energy associated with it Electrons are confined to specific

energy levels In order to move from one level to the

next, an electron must absorb or release a certain quantum of energy

The Bohr Model and Electron Transitions Illustrated

The color of the light you see is related to the amount of energy being released!!

A QUANTUM of energy is absorbed to jump levels from GROUND STATE to EXCITED STATE

Returning to GROUND STATE , the electron releases energy as PHOTON(S) of light!

Operation of Spectroscope

The Bohr Model

Let’s draw some Bohr diagrams!!!!

• 11Na

• 1H

• 2He

IV. The Quantum Mechanical Model of the Atom

Erwin Schrodinger, 1887-1961, Austrian“Hmmm…Matter also has particle

and wave characteristics. So

since matter is made of atoms,

and atoms contain electrons, then electrons could also travel in

waves!! THAT’S IT!!! I’ll be

famous!! Won’t Mom be proud!!”

Proposed the quantum mechanical model for electrons.

The exact path of the electron can not be determined because it is traveling near the speed of light and is too small in size.

This idea was based on the work of Werner Heisenberg, 1901-1976, German.

In the Heisenberg Uncertainty Principle there is a limit to how certain we can be about the position and speed of very tiny particles such as electrons.

Werner Heisenberg, 1901-1976, German

“Its impossible to know the

position and the speed of an

electron at any given moment—

kind of like trying to see

Road Runner’s legs when he’s running from

Wylie Coyote!”

Heisenberg Uncertainty PrincipleWhere’d it go?? Where’d it go???

In the quantum model only the probability of finding the electron in a certain area can be determined.

The most highly probable location for an electron about the nucleus is the orbital.

The combination of these areas about the nucleus is called the electron cloud.

Objectives #8-10: Quantum Numbers

I. Schrodinger’s Equation:EΨ = -h2/2m(ð2Ψ/ðx2 + ðΨ/ðy2 + ð2Ψ/ðz2) + V(x, y, z)Ψ (Neat, huh? No you do NOT have to use it,

memorize it or solve it!!! You’re welcome! But it IS cool… )

Solving the previous equation produces various orbital shapes, just as solving

y = 1/2x + 2 produces a straight line.

II. Quantum Numbers Describe energy and location of

electrons Every electron in an atom is

unique; each electron has a different energy and therefore will have a different set of quantum numbers.

Evolution of the Bohr Model into the Quantum Mechanical Model

The energy level number is equal to the number of subshells within that energy level…

A. Principle Quantum Number (n) Indicates energy and distance from nucleus Indicates energy level number Can take on values of: 1 infinity, but 1 7 is currently verified & understood.B. Orbital (Angular Momentum) Quantum

Number (l) Indicates shape of orbital (sublevel) Can take on values of: 0 n-1 (0,1,2,3,etc) Orbitals that have the same value of n, but

different l values are in sublevels, which are designated by letters to avoid confusion!

Possible Orbital Shapes:

“l” value Shape Letter Designatio

n0 Sphere s1 Dumbbell

(pronounced)P

2 Double dumbbell(diffuse)

d

3 Multi-lobed f4* *Predicted, but

not verifiedg

Orbital Shapess = 1 direction

p = 3 directions

d = 5 directions

f = 7 directions

The number of orbital (sublevel) shapes in a level is equal to the level number:

Level Number

# of Shapes Allowed

Shapes

1 1 s2 2 s p3 3 s p d4 4 s p d f

C. Magnetic Quantum Number (ml) Indicates the orientation

(direction) of the orbital in space Indicates the number of orbital

directions in a sublevel Can take on values of: –l 0 +l

C. Magnetic Quantum Numbers (ml)

“l” value Sublevel Designatio

n

“ml” Values Allowed

# of Orbitals

0 s 0 11 p -1, 0, +1 32 d -2, -1, 0, +1,

+25

3 f -3, -2, -1, 0, +1, +2, +3

7

D.Spin Quantum Number (ms) Indicates the direction of electron

spin Can take on values of: +1/2, -1/2 No more than 2 electrons can occupy

a single orbital

III.Problems Involving Quantum Numbers

See notes

IV. Summary of Electron Energy Level Capacities

(See Chart in Lecture Guide) Some relationships to notice:

If “n” is the number of levels, then the number of sublevels is equal to “n”If “n” is the number of levels, then the total number of orbitals in a level is equal to n2

If “n” is the number of levels (and every orbital can hold up to 2 electrons), then the total number of electrons in a level is equal to 2n2

Objectives #11-12: Electron Configurations

Electron configurations show electron arrangement

I. Rules Governing Electron Configurations1. The Aufbau Principle

Electrons enter orbitals of lowest energy

first2. The Pauli Exclusionary Principle

An atomic orbital may describe, at most,

2 electrons

3. Hund’s RuleElectrons enter orbitals of the same energy with the same

spin until each orbital contains one electron before pairing begins

II. Examples of Electron Configurations w/Orbital Notations & Noble-gas Configurations

1H:

2He:

3Li:

6C:

15P:

The Diagonal RuleUsed to track the order of electrons as they fill available orbitals, according to the Aufbau Principle. Remember that all electrons fill using the LOWEST amount of energy possible!

More Examples…Use the Diagonal Rule!

19K:___________________________

26Fe:__________________________

31Ga:__________________________

38Sr:_________________________

54Xe:_________________________

Noble Gas Configurations:Write the symbol of the closest noble gas that is LOWER in atomic number. Then, write out the remaining part of the configuration, following the diagonal rule.

Example: 74W—Tungsten

Exceptions to the Aufbau Principle:24Cr:

29Cu:

Objectives #13-20: The Periodic Table & Periodicity of Properties

I. Development of the Periodic Table The work of Mendeleev (1871,

Russian)Elements were grouped by their

properties; Allowed for prediction of new

elements; Elements arranged by increasing

atomic mass

Dimitri Mendeleev,1834-1907, Russian

“I ALMOST have it…if I could just find those

pesky missing pieces!!”

Dimitri Mendeleev’s Periodic Table - 1871

Henry Moseley,1887-1915,English

“Hmmm…what if we

used atomic NUMBERS rather than

mass…”

Objectives #13-20: The Periodic Table & Periodicity of Properties

The work of Moseley (1911, English): Used X-ray studies to determine

atomic numbers of elements;Arranged in order of increasing

atomic number

Periodic Table of 1930 Based on Henry Moseley’s Work

Glenn Seaborg, 1912-1999, American

“Now, if I move this here, and slide this

down over there…”

The work of Seaborg (1951, American)Contributed to the discovery of 10 elements:

plutonium, americium, curium, berkelium, californium, einsteinium, fermium, mendelevium, nobelium and element 106, which was named seaborgium in his honor while he was still living.

Pioneer in nuclear medicine and developed numerous isotopes of elements with important applications in the diagnosis and treatment of diseases, most notably iodine-131, which is used in the treatment of thyroid disease

Developed the “Actinide Concept”, which led to a reorganization of the PT

Modern Periodic Table Based on the Work of Glenn Seaborg, 1951

II. Organization of Modern Periodic Table

Groups Periods Metals Nonmetals Metalloids Representative elements

Alkali MetalsAlkali Earth MetalsHalogensNoble Gases

Transition Metals Rare Earth Metals

LanthanidesActinides

III. Valence Electrons and the Periodic Table

In most chemical reactions only the valence electrons are involved

Valence electrons are the outer level electrons in an atom

Examples—see lecture guide

Valence Electrons Illustrated

Objectives #13-20: The Periodic Table / Periodicity of Properties

IV.Periodic Properties Periodic Law: The properties of the elements vary

periodically—in a predictable pattern— when placed in order of increasing atomic

numbers Periodic Trends

1. Atomic Radii1/2 the distance between the nuclei of

identical atoms that are bonded together (how big the atom is!)

Objectives #13-20: The Periodic Table / Periodicity of Properties

Increases down a groupDecreases across the period

A cation or positive ion is smaller than its original parent atomAn anion or negative ion is larger than its original parent atom

2. Ionization EnergyThe amount of energy required to remove an electron from an atom

Objectives #13-20: The Periodic Table / Periodicity of Properties

decreases down a groupincreases across the period

3. Electronegativity A measure of the ability of an atom

to gain electrons during the bonding process

decreases down a group increases across the period

Objectives #13-20: The Periodic Table / Periodicity of Properties

Why these variations occur: Adding additional energy levels

increases size of atom The “shielding effect”--inner energy

levels block influence of nucleus from outer energy levels--also increases

Increasing nuclear charge (more protons) holds electrons slightly closer to the nucleus across period