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 Erwin Schrodinger & his math equations
Bohr Model 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 (aka the versions youve learned before) Either gain or lose energy to move levels Electrons CANNOT be between levels 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 Has properties of both: a Particle ( ____________) a Wave Light Waves: Amplitude: crest of the wave (height from 0) Wavelength: distance between crests ()) Frequency: # of waves per unit time ()) Units: Hertz (Hz) aka s-1 Math Time!!!
c = C = speed of light (constant) = 2.998 x 10 8 m/s = Wavelength (m) = Frequency (Hz or s-1) More Math The energy (E) of a photon is directly proportional to its frequency. Higher freq = More Energy Lower Freq = Less Energy E=hxv
E = energy (joules J) H = Planks constant = 6.626E-34 J/s v = Frequency (Hz or s-1) Example: What is the energy of a quantum of light with a frequency of 7.39 x 1014 Hz? Think about this E =hxv c = 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 Look at the gas tubes and follow directions provided. Continuous Spectrum v. Line Spectrum What Lab? did you observe in the Emission
Light has Wave-Particle Duality (& so do electrons) Particle Depends on experiment / what we try to observe Throws & Wave-like Nature 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 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
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
3d_____ 4s _____ _____ _____ ____ _____ 3p ______ ______ ______ 3s ______ 2p ______ ______ ______ 2s ______ 1s ______ Butall of these electrons are not organized into hotel rooms, but ATOMIC ORBITALS
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 Quantum Numbers Get readyhere we go 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 (thinkEnergy 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
Angular Momentum Quantum Number (Azimuthal Quantum Number) (thinkenergy sublevel) Pay attentionthis is where it starts to get complicated l Allowed values: 0, 1, 2, (n-1)
Describes the sublevel SHAPE of the orbital SHAPES: l l
l l = = = = 0 1 2 3 = =
= = s orbital = spherical cloud p orbital = dumbbell cloud d orbital = clover cloud f orbital = too complicated Example If I had a principal quantum number of 2, what are my possible angular momentum quantum numbers? n=2
l= Angular Momentum Quantum Number: Orbital Shapes Magnetic Quantum Number (ml) Determines Possible spatial orientation (x, y, z, plane) Values: - l to + l
Examples: if it is a d orbital d orbital: l= ml = 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 ml = 0 Energy level = Sublevel = _____ - orbital Orientation: Orbital: Number of orbitals within an energy level: n 2 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 ms Describes the direction of the electrons spin within
an orbital (remember, each orbital only holds 2 electrons) Possible Values: or - (spin up, spin down) Think back to hogs Ahhhits too much informationHELP!!! Solution: STUDY and PRACTICE!!!
Quantum # Symbol Possible Values Description Principle Quantum Number n 1, 2, 3, etc
Energy level Angular Momentum Quantum Number l 0 n-1 Sublevel & shape Magnetic Quantum Number
ml -l +l Spatial Orientation of orbital (x,y,z) Spin Quantum Number ms + or -
Direction of Spin Examples 1. n = 3 (what are the possible quantum numbers?) 2. What orbital corresponds to n = 4 & l = 2? What
orbital corresponds to n = 4 , l = 1, ml = -1 Energy Level = Sublevel = Orbital orientation = Orbital = Re-iterate: Orbital How Many Types of Orbitals (orientations)
s 1 p 3 d 5 f
7 How Many Electrons in Shape Principle Quantum Number (n) 1 2 3 4 5
6 7 Angular Momentum Quantum Number (sublevels) (l) Shapes of Sublevels # electrons (2n2)
Principle Angular Quantum Number Momentum (n) Quantum Number (sublevels) (l) Shapes of Sublevels # electrons (2n2)
1 0 s 2 2 0, 1 sp 8
3 0, 1, 2 spd 18 4 0, 1, 2, 3 spdf
32 5 0, 1, 2, 3, 4 s p d f (g) 50 6 0, 1, 2, 3, 4, 5 s p d f (g h)
72 7 0, 1, 2, 3, 4, 5, 6 s p d f (g h i) 98 STOP Do You Have Any Questions?
PART 3 Rules of Electron Configuration 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 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? Hunds 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 Writing Electron Configurations
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
2s ___ 1s ___ Review: 1. How many electrons fill an s orbital? 2. How many electrons fill a p orbital ? (remember subshells) 3.
How many electrons fill a d orbital? 4. How many electrons fill an f orbital? Example: Cl 3d ___ ___ ___ ___ ___ 4s ___ 3p ___ ___ ___ 3s ___ 2p ___ ___ ___ 2s ___ 1s ___ Give the final E.C:
With a partner: Examples: Give the E.C H He Li Be B C N F No moreMake it stop! @!!!! Write
the electron configuration for Barium: Ahhhhhhhhhh!!! But waittheres a shortcut Noble Too many electrons!!
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 Examples: Sb Stop & Practice E.C. Worksheet
All Together Now Mendeleev didnt know quantum numbers BUTour periodic table is related to HOW electrons fill the levels in the different shells Blocks s block
Groups p Block Groups 1&2 38 d block Transition
Elements f Block Rare earth metals It ends w/ Another Example: Ba (shorthand) Stop & Practice
Patterns in Electron Configuration Worksheet 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
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