chemistry i honors--unit 3: quantum mechanical model of the atom & periodic trends
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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 Radiation - PowerPoint PPT PresentationTRANSCRIPT
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