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Quantum Mechanics & Electron Configuration. Chapter 5: Electrons in Atoms. Part 1: Models of the Atom. 1897: Thompson Model (Plum Pudding) 1911: Rutherford Model – Small, dense, + charged nucleus Electrons orbit around 1913: Bohr Model 1926: Quantum Mechanical Model –

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quantum mechanics electron configuration

Quantum Mechanics & Electron Configuration

Chapter 5: Electrons in Atoms

part 1 models of the atom
Part 1: Models of the Atom

1897: Thompson Model (Plum Pudding)

1911: Rutherford Model –

Small, dense, + charged nucleus

Electrons orbit around

1913: Bohr Model

1926: Quantum Mechanical Model –

Erwin Schrodinger & his math equations

bohr model aka the versions you ve learned before
Bohr Model (aka the versions you’ve learned before)
  • Electrons move around the nucleus in fixed spherical orbits with fixed energies
    • Fixed energies = orbits / energy levels
      • Aka rungs of a ladder
  • Electrons can go to a higher or lower energy level
    • Either gain or lose energy to move levels
  • Electrons CANNOT be between levels
atomic emission spectra
Atomic Emission Spectra

** When atoms absorb energy (i.e. electric current), they move to a higher energy level …

… these electrons emit light when they return back to a lower energy level

  • Emission spectra is unique for each element
    • The light emitted consists of only a mixture to specific frequencies…
  • If you pass the light through a slit and then a prism, you can separate the resulting light into its frequencies (aka colors)

Barium

light
Light
  • Has properties of both:

 a Particle ( ____________)

 a Wave

Light Waves:

Amplitude: crest of the wave (height from 0)

Wavelength: distance between crests (λ)

Frequency: # of waves per unit time (ν)

Units: Hertz (Hz) aka s-1

math time
Math Time!!!

c = λν

C = speed of light (constant) = 2.998 x 108 m/s

λ = Wavelength (m)

ν = Frequency (Hz or s-1)

more math
More Math…
  • The energy (E) of a photon is directly proportional to its frequency.

Higher freq = More Energy

Lower Freq = Less Energy

E = h x v

E = energy (joules – J)

H = Plank’s constant = 6.626E-34 J/s

v = Frequency (Hz or s-1)

example
Example:

What is the energy of a quantum of light with a frequency of 7.39 x 1014 Hz?

think about this
Think about this…
  • E = h x v
  • c = λν

What would you do if you were asked to solve for the frequency of light if you are given a wavelength of 700nm?

What would you do if you were asked to find the energy of light if you are given a wavelength of 480nm?

emission spectra lab
Emission Spectra Lab

Look at the gas tubes and follow directions provided.

continuous spectrum v line spectrum
Continuous Spectrum v. Line Spectrum
  • What did you observe in the Emission Lab?
light has wave particle duality so do electrons
Light has Wave-Particle Duality (& so do electrons)
  • Particle & Wave-like Nature
    • Depends on experiment / what we try to observe
  • Throws a wrench in Bohr Model…
    • New method of describing the motion of subatomic particles

= foundation of quantum mechanics = movement/organization of subatomic particles

the quantum mechanical model
The Quantum Mechanical Model
  • This is what we use today
  • Describes: LOCATION & ENERGY of electrons
    • Electrons do not have a direct orbit around nucleus
      • Based on probability
      • Electron clouds
    • Electrons do have energy levels
hog hilton sample problem
Hog Hilton Sample Problem
  • Book 15 hogs into their rooms
    • 6th floor ____ ____ ____ _____ _____
    • 6th floor ______
    • 5th floor ______ ______ ______
    • 4th floor ______
    • 3rd floor ______ ______ ______
    • 2nd floor ______
    • 1st floor ______
hog hilton sample problem place 15 electrons into their spaces
Hog Hilton Sample ProblemPlace 15 electrons into their spaces
  • 3d_____ _____ _____ _____ ____
  • 4s _____
  • 3p ______ ______ ______
  • 3s ______
  • 2p ______ ______ ______
  • 2s ______
  • 1s ______
so what exactly is an atomic orbital
So, what exactly is an ATOMIC ORBITAL?

Atomic Orbital = region of space in which there is a high probability of finding an electron

  • They come in different SHAPES, SIZES & ENERGY LEVELS!!
  • These are described by Quantum Numbers…
part 2

Part 2

Quantum Numbers

Get ready…here we go…

quantum numbers
Quantum Numbers

Used to describe the location of electrons

Electrons in an atom CANNOT have the same quantum numbers

 Unique for each electron

 Like an address

principle quantum number think energy level
Principle Quantum Number (think…Energy Level)
  • n
  • Allowable values = 1, 2, 3 … n (positive, integer values)
  • Describes energy level
    • Position of the electron w/ respect to nucleus
    • As n increases = further from nucleus
slide23
Angular Momentum Quantum Number(Azimuthal Quantum Number)(think…energy sublevel)Pay attention…this is where it starts to get complicated
  • l
  • Allowed values: 0, 1, 2, … (n-1)
  • Describes the sublevel
    • SHAPE of the orbital
  • SHAPES:
    • l = 0 = s orbital = spherical cloud
    • l = 1 = p orbital = dumbbell cloud
    • l = 2 = d orbital = clover cloud
    • l = 3 = f orbital = … too complicated
example1
Example
  • If I had a principal quantum number of 2, what are my possible angular momentum quantum numbers?

n = 2

l =

magnetic quantum number m l
Magnetic Quantum Number (ml)
  • Determines spatial orientation (x, y, z, plane)
  • Possible Values: - l to + l
    • Examples: if it is a d orbital

d orbital:

l =

ml =

example p orbital
Example: p-orbital

n = 2

l =

ml =

This means, there are _______ p-orbitals and that they are in three directions (x, y, z axes):

what orbital corresponds to n 2 l 1 m l 0
What orbital corresponds to :n = 2l = 1ml = 0

Energy level =

Sublevel = _____ - orbital

Orientation:

Orbital:

slide29

Number of orbitals within an energy level: n2

Examples: How many orbitals are in energy level 2?

n =

l =

ml =

Orbitals =

  • Each orbital holds 2 electrons:So, how many electrons can energy level 2 hold?

# Electrons = 2n2

spin quantum number
Spin Quantum Number
  • ms
    • Describes the direction of the electrons spin within an orbital (remember, each orbital only holds 2 electrons)
    • Possible Values: ½ or -½ (spin up, spin down)
      • Think back to hogs…
ahhh it s too much information help
Ahhh…it’s too much information…HELP!!!
  • Solution: STUDY and PRACTICE!!!
examples
Examples
  • n = 3 (what are the possible quantum numbers?)
  • What orbital corresponds to n = 4 & l = 2?
slide33

What orbital corresponds to

n = 4 , l = 1, ml = -1

Energy Level =

Sublevel =

Orbital orientation =

Orbital =

part 3
PART 3

Rules of Electron Configuration

aufbau principle
Aufbau Principle
  • Electrons enter orbitals of lowest energy first
    • Orbitals within a sublevel have equal energy

(3px, 3py, 3pz)

    • Exceptions: Cr , Cu
  • Which hog rules is this?
pauli exclusion principle
Pauli Exclusion Principle
  • An atomic orbital may only hold two electrons
    • Electrons must have opposite spin
      • Clockwise or counterclockwise spin
      • Denoted with arrows
      • Prevents two electrons from having same quantum numbers
  • Which hog rule is this?
hund s rule
Hund’s Rule
  • Every orbital of the same energy is singly occupied before any orbital is doubly occupied
    • Electrons have the same spin
    • Second electrons added have opposite spins
  • Which hog rule is this?
part 4
PART 4

Writing Electron Configurations

electron configuration diagonal rule
Electron Configuration Diagonal Rule
  • Starting with the top arrow, follow the arrows one by one in the direction they point, listing the sublevels as you pass through them.
  • Stop when you get to the sublevel you need.
electron orbital diagram
Electron Orbital Diagram

3d ___ ___ ___ ___ ___

4s ___

3p ___ ___ ___

3s ___

2p ___ ___ ___

2s ___

1s ___

example fill orbitals w 7 electrons
Example: Fill Orbitals w/ 7 electrons

3d ___ ___ ___ ___ ___

4s ___

3p ___ ___ ___

3s ___

2p ___ ___ ___

2s ___

1s ___

review
Review:
  • How many electrons fill an s orbital?
  • How many electrons fill a p orbital ?(remember subshells…)
  • How many electrons fill a d orbital?
  • How many electrons fill an f orbital?
example cl
Example: Cl

3d ___ ___ ___ ___ ___

4s ___

3p ___ ___ ___

3s ___

2p ___ ___ ___

2s ___

1s ___

Give the final E.C:

no more make it stop @
No more…Make it stop!@!!!!
  • Write the electron configuration for Barium:
  • Ahhhhhhhhhh!!! Too many electrons!!
  • But wait…there’s a shortcut…
  • Noble gas / shorthand configuration:
    • Find the nearest noble gas that came before the element you are interested in
    • Write the symbol of that noble gas in [brackets]
    • Write the configuration as normal from there…
stop practice
Stop & Practice
  • E.C. Worksheet
all together now
All Together Now…
  • Mendeleev didn’t know quantum numbers
    • BUT…our periodic table is related to HOW electrons fill the levels in the different shells
  • Blocks
    • s block
      • Groups 1 & 2
    • p Block
      • Groups 3 – 8
    • d block
      • Transition Elements
    • f Block
      • Rare earth metals
stop practice1
Stop & Practice
  • Patterns in Electron Configuration Worksheet
columns
Columns
  • Elements have similar properties
    • Why?
      • Similar ground state electron configurations
    • Examples
      • Noble gases
        • Complete sublevel
          • Favorable - do not react
      • Halogens
        • One electron short of completely filled sublevel
          • Readily react with elements who have a single electron