Download Electrons BellwoodNotes

Electrons in an Atom
• According to the Bohr Model,
electrons (e-) can only orbit
the nucleus in specific, allowed
pathways.
• They move toward and away from the
nucleus by “steps” or discrete amounts of
energy that are released or absorbed.
• e- farther from the nucleus have more energy.
Those closer to the nucleus have less energy.
The Problem (s)…
• If electrons orbited the nucleus,
they would constantly be accelerating
• Accelerating charges give off all
colors of light at the same time
• White light is made up of every visible color of
light
• There’s nothing stopping the electrons from
crashing into the nucleus? (+ attract -)
Solution: Very similar to a ladder. Just as you cannot
step on the air between the rungs, an electron
cannot exist between the levels.
Light is going to help us understand
the atom!
We need to know…
• What is light?
• How does light behave?
• What does light have to do with electrons?
• Can we explain/predict the kind
of light that is given off by different
elements?
The Electromagnetic Spectrum
All the types of light!
ROY G. BIV
Think… Slinky
Light is a Wave
Light – as a Wave (How do we know this?)
 Can interfere with itself (two waves crash together
and cancel each other out)

Can refract (be bent as it passes through different
materials)
Parts of a Wave
Light – as a Wave
 Wavelength (λ – “lambda”): the distance between
the peaks

Frequency (ν – “nu”): number of waves that pass a
given point in a specific amount of time
λ
crest
trough
Light as a wave
speed of light = frequency x wavelength
c=νλ
(where c is always equal to 3.00 x 108 m/sec)


Longer wavelength → _________ frequency
Shorter wavelength → _________frequency
In the visible spectrum, color is associated with
different wavelengths (ROY G. BIV)
Longer λ and lower ν: towards Red
Shorter λ and higher ν: towards Violet
Einstein’s Theory
To make sense of the fact that light can eject
electrons even if its intensity (brightness) is low…

Light is both waves and particles –think
“stream of particles”

Particle: The photon – it has zero mass and
carries a quantum (piece) of energy

If the energy of the light particle is too low, no
matter how many particles hit the metal, no
electrons will pop off!
New Ideas about Light
Prior to the 1900s – light was thought of as only a wave
… Now there is a duality of light
 “Duality” means that there are two ways to represent it
Wave

Particle
Properties of both are present
This also means that… all particles have a wave nature, all
waves have a particle nature.
Turn in to me:
1. Rank the colors of light from low to high
energy.
2. Rank the colors of light from small to large
wavelength
Violet
Blue Green Yellow Orange Red
Particle of Light



Emission of light by hot objects happens in small specific
amounts (packets)– quanta
Each particle carries with it an amount of energy which
depends on the frequency of light
Quantum – that small bundle of energy
Energy = Planck’s Constant * frequency emitted
E=hν
(where h = 6.626 x 10-34 J*s)
Spectroscopic analysis of the visible
spectrum…
…produces all of the colors in a continuous spectrum
Spectroscopic analysis of the hydrogen
spectrum…
…produces a “bright line” spectrum

Ground State – lowest energy state


This means that e- are found in shells closer to the nucleus.
Excited State – higher potential energy of an atom


A form of heat, light, electrical, or mechanical energy is needed
to go from the ground to an excited state
As electrons increase in energy, they move away from the
nucleus and into outer shells.
Atoms and Light

Absorption (take in) of energy moves electrons
from a ground state to a higher energy state


Heat, light, electrical, mechanical energy
Emission (or give off) of energy lets electrons fall
back down to a lower energy state


Always light
We don’t know why this happens
Excited State
Energy
Going In
Ground State
Energy Coming
Out
“Color”
The kind of light given off (emitted) by each element is
UNIQUE because each element has a UNIQUE arrangement
of electrons! (Think fingerprint)
(It’s not just about electrons in “shells”… it’s more
complicated than that!
Hydrogen’s Line Emission Spectrum
Excited State
Energy (Light)
Ground State
This led to the quantum model of the atom!
Electrons are in orbitals, not orbits
Not

Come from a mathematical equation
Orbitals (3-D Fuzzy Shapes)




Three dimensional space that electrons most probably occupy
The math equation treats electrons like waves
You can solve the equation to get the shape of space in which
electrons are
Shapes look like “clouds” of probability
What is the “address” of an electron?

Where do you live?
- Where on the block
- Style of house
- Which way is your house facing

Electrons are identified the same way …
- Principle Energy Level (identified by n)
- Sublevel (identified by l)
- Orbital Orientation (identified by ml)
- Spin State (identified by ms)
Four Quantum Numbers…
1. Principle Quantum Number (n )
• Indicates the main energy level occupied by the e- (distance
from the nucleus)
•
Shell Number (1st shell is closest to nucleus, 2nd is further, and
so on …)
•
Come from the Bohr Model
•
Values of n can only be
positive integers (1, 2,
3, etc.)
Four Quantum Numbers…cont.
2. Angular Momentum Quantum Number ( l )
• Indicates the shape of the orbitals (orbital = individual 3D
shape in space where the probability of finding e- is
greatest)
•
Nickname is subshell of n
•
Designated s, p, d, f
•
Values of l are zero and all positive integers less than
equal to n - 1
The number of subshells in a shell is equal to the
shell number



1st shell – 1 subshell
2nd shell – 2 subshells
3rd shell – 3 subshells
s subshell
 Spherical shaped
 One orbital, 2 e-
p subshell
 Dumbbell shaped
 Three orbitals, 6 ed subshell
 Double peanut shaped
 Five orbitals, 10 ef subshell
 Flower shaped
 Seven orbitals, 14 e-
What are the two quantum numbers
learned so far?
1. Principal Quantum Number
shell number 1-7
distance from nucleus
2. Angular Momentum Quantum Number
subshell s p d f
actual space where the probability of finding
an e- pair is greatest (orbital)
Four Quantum Numbers…cont.
3. Magnetic Quantum Number (ml )
• Indicates the orientation of an orbital (the plane that the
orbital is in)
• Because an s orbital is spherical, it only has one orientation
(ml = 0)
• p orbitals can have three different orientations, one along
the x-axis, one along the y-axis, and one along the z-axis (ml
= -1, ml = 0, ml = +1)
• (just an illustration: I don’t need you to memorize)
4. Spin Quantum Number (




s
Indicates the two spin states of an e- in an orbital
Only 2 e- fit in each orbital, and they spin in opposite
directions (up and down!)
Possible ms values are - ½ , + ½
Spin is represented by dashes inside circles
Empty

m)
Half-Filled
Filled
Half-Filled
Filled
or arrows
Empty
What are the four quantum numbers?
Principal Quantum Number
shell number 1-7
distance from nucleus
2. Angular Momentum Quantum Number
subshell s p d f
actual space where the probability of finding an
e- pair is greatest (orbital)
3. Magnetic Quantum Number
the # of directions your shape (orbitals) faces
4. Spin Quantum Number
clockwise, counterclockwise
1.
Bohr Models
Represented with pictures…
Mathematical Models
Represented with… ?
Summary Table
Shells Subshell Orbitals
1
2
3
4
s
1
s
p
s
1
3
1
p
d
s
3
5
1
p
d
3
5
f
7
Total #
orbitals
in shell
1
4
9
16
#e-
Total #ein shell
2
2
2
6
2
6
10
2
6
10
14
8
18
32
It is a bit complicated…
shells don’t always get filled from 1 to 2 to 3
etc. because some subshells overlap
Start End
Nucleus 1
Start
Start
2
End
3
End Start
Start
End
5
4
End
What order are they filled in?
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
7p
3d
4d
5d
6d
?
4f
5f
?
?
?
?
?
?
?
Electron Configurations
Orbital Notations
The Arrangement of electrons in
atoms
Aufbau Principle





“building up”
An electron occupies the lowest energy possible
The levels follow a pattern of increasing energy
Fill starting at nucleus (Bohr Models!)
p subshell  3 orbitals

Fill left to right
not
Hund’s Rule

Electrons do not pair up until there are no more
empty orbitals in that subshell
2e3e-
NOT
4e-
OR
Pauli Exclusion Principle

No two electrons in the same atom
will have the same set of four
quantum numbers
 No two electrons have the same
address

An orbital can only hold two electrons
because there are only two different
spin states (and the must be opposite)
To represent e- configurations, you need to know:
1. The atomic # (# of e-)
2. Order in which the e- fill
3. The number of orbitals in each sublevel (s, p, d, f)
(remember – they want a private room)
What to write:
1. Shell # (1, 2, 3, 4, …)
2. Sublevel letter (s, p, d, f)
3. # of e- orbital notation
(Remember how many e- are in each sublevel)
Electron configuration and the
periodic table
What do elements in column 1 have in
common?
Column 2?
Column 3?
What do elements in each of the rows have in
common?
Periodic Table

Different blocks on
the periodic table
(shaded in different
colors here)
correspond to
different types of
orbitals.
The lazy electron configuration

Start with the previous noble gas in brackets, and
then finish the rest of the configuration

“Condensed” electron configuration



Long way is Sodium = 1s2 2s2 2p6 3s1
Condensed Sodium = [Ne] 3s1
“valence” shell electrons are the ones listed as
following the noble gas (aka the outside shell
electrons)
Element
Lithium
Configuration
notation
1s22s1
[He]2s1
____
1s
Beryllium
____
____
2p
____
____
2s
____
____
2p
____
[He]2s2p2
____
2s
____
____
2p
____
1s22s2p3
[He]2s2p3
____
2s
____
____
2p
____
1s22s2p4
[He]2s2p4
____
2s
____
____
2p
____
1s22s2p5
[He]2s2p5
____
1s
Neon
____
2s
1s22s2p2
____
1s
Fluorine
____
[He]2s2p1
____
1s
Oxygen
____
2p
1s22s2p1
____
1s
Nitrogen
____
[He]2s2
____
1s
Carbon
____
2s
1s22s2
____
1s
Boron
Noble gas
notation
Orbital notation
____
2s
____
____
2p
____
1s22s2p6
[He]2s2p6
____
1s
____
2s
____
____
2p
____
Valence Shell Electrons



Valence = outermost
Count the total electrons in the highest shell
number
Do not count electrons in d subshells
How many valence shell electrons are in the
following elements?





H
He
O
S
Br