Wave Particle Duality & Electron Configurations
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Transcript of Wave Particle Duality & Electron Configurations
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Wave Particle Duality &Electron Configurations
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Remember Rutherford? Proposed model of the atom had a nucleus of
positive charge surrounded by a relatively large area of empty space where electrons orbited
Did not propose an arrangement for the electrons
Did not explain why the electrons were not pulled into the nucleus (attraction of opposite charges)
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Light as a Wave Light exhibits characteristics of waves
Wavelength (λ)
Frequency (ν)
amplitude
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Electromagnetic Waves Different types
Each type has characteristic λ and ν
All parts travel with the speed of light (c)… 3.00 x 108 m/s
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Electromagnetic Wave Formula
𝑐= λν
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EXAMPLE PROBLEM 5.1 Microwaves are used to cook food and
transmit information. What is the wavelength of a microwave that has a frequency of 3.44 x 109 Hz?
Now You Try:Practice Problems 1 – 4 on page 140
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The Particle Nature of Light Objects that are heated often give off a
characteristic color (red of stove burner, white of light bulb)
View of light as a wave did not provide an accurate explanation of why this occurs
So….
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Max Planck (1858 – 1947) Concluded that energy could only be
gained or lost in small, specific amounts (like tiny packages)…called these amounts quanta
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Energy of a Quantum
E = hν
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The Photoelectric Effect Another phenomenon that could not be
explained with light as a wave
When light of a certain minimum frequency shines on a metal’s surface, the metal will eject electrons (video)
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Example 5.2 Every object gets its color by reflecting a
certain portion of incident light. The color is determined by the wavelength of the reflected photons, thus by their energy. What is the energy of a photon from the violet portion of the Sun’s light if it has a frequency of 7.230 x 1014 s-1?
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Atomic Emission Spectra Also called line spectra…not continuous
Set of frequencies of electromagnetic waves emitted by an element
Not continuous
Unique for each element (like a fingerprint)
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Electron Configurations Bohr
› Model stated that atoms orbit the nucleus in definite paths (energy levels)
› Patterned this model after planets orbiting the sun
› Electrons in a particular path (energy level) have a fixed amount of energy…quantized
› Energy levels are analogous to the rungs on a ladder
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Bohr’s Model Based on a hydrogen atom
Assigned a quantum number (n) to each orbit
As value of n increases, the amount of energy increases
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Energy relationships…
∆𝐸=𝐸 (h h𝑖𝑔 𝑒𝑟 𝑜𝑟𝑏𝑖𝑡 )−𝐸 (𝑙𝑜𝑤𝑒𝑟 𝑜𝑟𝑏𝑖𝑡 )=𝐸 ( h𝑝 𝑜𝑡𝑜𝑛 )=h𝑣
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Limitations of Bohr’s Model Worked well for hydrogen
Did not explain the atomic spectra produced by any other elements
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The Quantum Mechanical Model of the Atom
Louis de Broglie (1892 – 1987) › Recognized that electrons exhibited
characteristics of waves› Recognized that light has properties of
both waves and particles› Theorized that matter must be able to
possess qualities of waves and particles as well
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de Broglie Equation: Predicts that all matter has wave
characteristics
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Heisenberg Uncertainty Principle
States that it is fundamentally impossible to know both the precise velocity (momentum) and location of a particle at the same time
Applies to all matter, but is useful only with really small particles…like electrons
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Electron Configurations Schrodinger
› Revised Bohr’s model› Mathematical equation to determine the
most likely place an electron would orbit the nucleus
› Gives the probability of finding an electron in a particular place within the atom
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Quantum Numbers Used to describe orbitals Specify the properties of atomic
orbitals and the properties of the electrons in the orbitals
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4 Quantum Numbers First three derived from the
Schrodinger equation:› Main energy level (n)› Shape of orbital (l)› Orientation of orbital (ml)
Fourth is the spin quantum number (ms)› Describes the fundamental state of the
electron
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Principal Quantum number Symbolized by n Indicates the main energy level occupied
by an electron Values are positive integers As n increases, so does the distance from
the nucleus More than one electron can have the same
n Total number of orbitals for a given energy
level is given by n2
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Angular Momentum Quantum Number
Each main energy level (except the 1st) has different orbitals of different shapes
Symbolized by l Indicates the shape of the orbital The number of orbital shapes possible
for each energy level is equal to the value of n
The values of l allowed are zero through n-1
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Angular Momentum Quantum Number
Depending on the value of l, the orbital is assigned a letter
0=s 1=p 2=d 3=f
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Shapes of sub-orbitals s = spherical p = dumbbell shaped d= complex f = way to complicated to explain…see
illustrations
Atomic orbitals are designated by the n followed by the letter of the sublevel
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Magnetic Quantum Number Orbitals can have the same shape, but
different orientations around the nucleus
Magnetic quantum numbers indicate the orientation (ml)
s orbitals have only one orientation p orbitals can extend along the x, y, or
z axis 3 p sublevels (px, py, or pz)
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Magnetic Quantum Numbers
Values for m sublevels correspond values m = -1 m=0 and m=1
5 different d orbitals…therefore 5 different orientations
m=-2 m=-1 m=0 m=+1 m=+2 7 different f orbitals….7 orientations
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Spin Quantum Numbers Electrons spin on an internal axis Can spin in one of two possible
directions Spin quantum numbers can be +1/2 or
-1/2 A single orbital can hold a maximum of
2 electrons The electrons in a single orbital must
have opposite spins
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Electron Configurations Remember:
› All electrons can be described by a set of quantum numbers
› No two electrons can have the same set of quantum numbers
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Rules for Electron Configurations
Aufbau Principle› An electron will occupy the lowest energy
orbital that can receive it Pauli Exclusion Principle:
› No 2 electrons in the same atom can have the same set of quantum numbers
Hund’s Rule:› Orbitals of equal energy are each occupied by
one electron before any orbital is occupied by a second electron….All electrons in singly occupied orbitals have the same direction of spin (parallel spin)
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Energy Level Types of Orbitals
1 s
2 s, p
3 s, p, d
4 s, p, d, f
5 s, p, d, f
6 s, p
7 s, p
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Orbital Type # Sub-orbitals
s 1
p 3
d 5
f 7
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Sub-orbitals Each sub-orbital can hold two electrons
The electrons in a sub-orbital must have opposite spin
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Type of Orbital Max # electrons
s 2
p 6
d 10
f 14
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Energy Level Types of Orbital Max # of e-
1 s 2
2 s, p 8
3 s, p, d 18
4 s, p, d, f 32
5 s, p, d, f 32
6 s, p, 8
7 s, p 8
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Electron Configurations & the Periodic Table
s, p, d, f blocks
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s - block
1 s2 s3s4 s5s6 s7 s
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p - block
5p
2p3p4p
6p7p
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d - block
5d
3d4d
6d
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f - block
4f5f
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Give the Electron Configurations:
Carbon
Lithium
Sodium
Phosphorus
Neon
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You Try: Write the complete electron configuration
for:› Helium
› Sulfur
› Magnesium
› Silicon
› Tin
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Show orbital notations for: Carbon
Lithium
Sodium
Phosphorus
Neon
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You try Write the orbital notation for:
› Helium
› Sulfur
› Magnesium
› Silicon
› Tin
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Noble Gas Notations Use the noble gas that comes before
the element
Write the noble gas’s symbol in brackets
Continue with the electron configuration from there
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Noble Gas Notation Example Calcium:
› Electron Configuration:
› Noble Gas Notation:
You try these:a. Titaniumb. Siliconc. Barium