Unit 6 - Electron Structure (2).pptx

Unit 6: Electron Structures

and/or frequency of electromagnetic photons. 6.2: Calculate and interpret line spectra for identification of elements.

6.3: Compare the electron arrangements of historic electron models and the modern

quantum mechanical electron cloud models 6.4: Write and/or illustrate the bohr model,

electron configuration, or orbital diagram for an element using

Light Basics (electromagnetic

radiation)

• Light: Waves that oscillate in the electric and magnetic fields.

Wave properties

• _{Wavelength: Distance from}

crest-crest or trough-trough (m) • _{Amplitude: Height of wave from}

midpoint to crest

• Frequency: How many waves

pass through a space in a given time (Hz) (s-1)

• Period: Length of time between the passing of each wave (s)

Light Math

Main equations: λ • V = c and E = ħ • V

E = Energy of one photon of light (joules)

ħ = Planck’s Constant = 6.626 x 10-34 (J • s)

V = frequency (Hz) λ = wavelength (m)

c = Speed of light in a vacuum = 3.00 x 108 (m/s)

Or combine the equations for E = (ħ • c)/λ if needed…

Example

• _{Main equations: λ • V = c and E = ħ • V}

• _{Example: The sodium vapor lamp often seen on}

freeways in the US has a wavelength of 587.8 nm. Calculate the frequency of the light wave.

E = ħ • V = 6.626 x 10-34 (J • s) x 5.1 x 1014 (1/s) = 3.4 x

Now you try…



•

Calculate the energy of a 298 nm wavelength

light photon

•

Calculate the energy of a 3.4 moles of these

light photons

E = (ħ • c)/λ = (6.626 x 10

(J • s) • 3.00 x 10

m/s)

2.98 x 10

m

= 6.67047… x 10

J of 1 photon

6.022 x 10

photons • 6.67047… x 10

Photons

• _{Photon: The smallest particle of light that still retains the}

properties of light. (think “one single wavelength”)

• This is a way of treating light as a particle instead of waves

• _{Light photons are “quantized” in energy. Meaning, that each}

photon has a specific amount of energy.

• Niels Bohr found that the lines of a line spectra corresponded to a specific color, thus wavelength, thus frequency, thus

energy of light.

• _{He theorized and calculated that the color of photons emitted}

from an atom were proportional to the height of the “quantum leap” an electron undergoes!

• Basically: Specific size of quantum leap  1 color photon 

Bohr Model

• Bohr revolutionized the atomic model by theorizing that electrons “orbit” at specific distances from the nucleus. He called these specific orbits “Energy Levels.”

• _{Light is caused by electrons absorbing energy, becoming} “excited”, and jumping to higher energy levels. The

electrons could then “relax” and fall back to a lower

energy level, emitting a photon of a specific wavelength. • _{He could accurately calculate the wavelength of the light}

emitted by such an electron jump (called a “quantum leap”) using the Rydberg Equation: R_H = Rydberg

Constant =

1.096776 x 107

Atomic Emission

Energy

Level 1

Energy

Level 2

External

Energy

Nucleus

e-E1 Light wave

(photon) is

Now you try…



• _{What amount of energy would a photon of light}

contain that was emitted from a hydrogen atom as the result of a quantum leap and relaxation from the 4th

energy level to the 1st?

• _{RH = 1.096776 x 10}7 (1/m)

• _{h = 6.626 x 10}-34 (J • s)

• V = 1.028 x 107 Hz

• _{E = hv = 6.626 x 10}-34 (J • s) • 1.028 x 107 Hz = 6.813

Line Spectra and

Spectroscopy!

• _{Light from a source that passes through a prism is split}

up into a “spectrum” or a separation of light based on wavelength.

Spectroscopy

• Spectroscopy: Study of light spectra. (VERY powerful technique)

• _{If a specific compound, atom, element, gives off a}

specific line spectrum then you can tell its specific atomic composition!

• _{Can be used for infrared, x-ray, microwave, and}

Major Spectroscopy

Techniques

•

Radio

Nuclear Magnetic Resonance (NMR): Can ID Specific organic

molecules

Magnetic Resonance Imaging (MRI): Used to make a 3-D

image of the insides of organic matter

•

Infrared

IR Spectroscopy: Gives a specific “fingerprint” graph for each

compound.

•

Visible

Atomic Emission Spectra: Very basic/easy way to identify

elements

•

X-Ray

Quantum

Mechanics!

You guys already understand atomic theory, so this should be cake . BTW, after this you will understand what approximately 99% of Americans DON’T understand about how the world

fundamentally works.

Light: Wave or Particle?

• Photon: The smallest particle of light that still retains the properties of light. (think “one single wavelength”)

• _{This is a way of treating light as a particle instead of waves.} • So… We can brake waves up into particles if we treat each

wave like it is one particle…

• _{Can we treat particles like waves?}

• _{Are waves and particles the same thing?}

• _{What separates us from treating them as the same thing?} • What about a really small wave?

• _{What about a really fast particle?}

Wave-Particle Duality

•

Louis de Broglie asserted that if light waves could

be treated as particles then, basically, particles

(such as atoms) could be treated like waves!

•

He calculated that

ALL

matter and light

have

both

particle and wave properties.

matter=e-Uncertainty Principle

• _{Werner Heisenberg took things further… asserting that}

if ALL

matter had wave properties then ALL matter could not have a specific position in space/time like waves! Ever!

• Thus according to Heisenburg, there is a astronomically small

probability that you are not sitting in your chair right now but

are actually everywhere in this room at once… like a wave…

• Heisenberg Uncertainty Principle: The momentum

(direction of travel) and 3-D location of a small mass of matter (e-) cannot be known simultaneously with

Schrodinger Wave Function

•

Erwin Schrodinger solved the “psi equation”

for hydrogen for the possible locations of an

electron in the first energy level.

•

This generated the “electron cloud model”

for an atom.

No going back…

•

There is a 99.9% probability that an

electron will be found at one of the points

within these clouds.

Quantum Model

• _{Electron structure is incredibly important in chemistry!} • _{Electrons are responsible for how atoms bond, react,}

produce color, and most chemical/physical properties of substances!

• _{So, if electrons are really just statistical clouds of negative}

Quantum Mechanics and e- Cloud

Models

• This all finally lead to mapping out electron structures very specifically.

Electron Models from least to most specific:

• Bohr Model: Gives energy level for each electron.

• Electron Configuration: Gives energy level and sublevel for each electron.

• Orbital Configuration: Gives energy level, sublevel, orbital, and spin state for each electron.

• Quantum Numbers: Also gives energy level, sublevel,

Giving Each e- an “Address”

• _{Energy Level: A specific distance from the nucleus} (n=1,2,3,4,5…) (row #)

• Sublevel: A collection of orbitals of a certain shape at a certain energy level

(s, p, d, f) (blocks of periodic table)

• Orbital: A 3-D space where an e- is likely to be found (1/2 the rows in a block)

Rules for Orbital Filling

•

Aufbau Principle: e- enter orbitals of lowest

energy first

•

Pauli Exclusion Principle: An atomic orbital

contains a maximum of two electrons

•

Hund’s Rule: e- do not pair in an orbital (have

opposite spin) until all orbitals at that energy

level contain 1 e-

•

Octet Rule: The valence energy level can

Drawing Bohr Models

• _{Each row on the Periodic Table corresponds to an energy level.} • _{Draw 1 circle for each energy level.}

• _{Fill 1 electron for each element in each row.}

• _{Example: Draw the Bohr model diagram of elements from the}

following elements: Hydrogen, Nitrogen, and Aluminum

• The octet rule: An atom can have only 8 valence electrons MAX • All d-block element electrons get “stuffed” 1 level behind the

outermost energy level.

• _{All f-block electrons get “stuffed” 2 behind!}

Bohr Models

•

Draw a Bohr Model diagram of Sulfur

•

1 ring per energy level, 1 EL per row on

Periodic table

•

1 electron per element on table

•

Draw a Bohr Model diagram of Bromine

Now you try…



Sublevels and Electron

Configurations

• When the Schrodinger equation is solved for possible

electron locations, this results in a 99.9% probable area where electrons can be found.

• These statistical clouds take on very distinct sizes and shapes when the Schodinger equation is solved using different variables.

• 4 types of sublevels: s, p, d, and f. • _{s sublevels contain}

1 spherical orbitals • p sublevels contain

Sublevels and Electron

Configurations

• d sublevels contain 5 “x shaped

elliptical orbitals” on different axis

Drawing Electron

Configurations

• _{Electron Configurations: A “code” that gives a three} dimensional blueprint for the 3-D structure of an

atoms electrons.

• _{Adds more detail by adding} sublevel specifics

• I will show you a few examples: Hydrogen, lithium, iron, bismuth • _{I will translate bismuth from the}

“blueprint” to the 3-D picture as an example…

• _{Noble Gas Configurations:}

Start with “noble gas core” and continue. Easier to write 

Now you try…



•

What is the electron configuration of Carbon?

•

1s

2s

2p

•

What is the

noble gas

configuration for Gold?

•

[Xe]6s

4f

5d

•

What would a 3-D drawing of Sulfur look like based on

its e- config?

Aufbau Diagrams/Orbital

Configurations

• _{The aufbau diagram for Carbon and Vanadium}

• _{Similar to electron configurations but adds orbital and spin}

Now you try…



Quantum numbers!

• _{If you want to assign a UNIQUE and SPECIFIC address to each and}

every electron then you must use quantum numbers.

• _{AND you don’t want to draw a giant Aufbau diagram.}

• _{Quantum numbers specify a 4 number address for each electron.} • _{Each of the 4 numbers correlates to energy levels, sub levels,}

orbitals, and spin state

• _{Not grounds for testing on AP Test!}

• Principle quantum number (n) = energy level (row on periodic table)

• Angular quantum number (l) = sublevel (s,p,d or f)  (0,1,2,3)

• _{Magnetic quantum number (ml) = orbitals (s = 0) (p = -1,0,+1)}

(d=-2,-1,0,1,2)