I. Waves & Particles (p. 97-100) Ch. 4 - Electrons in Atoms.
I. Waves & Particles Ch. 6 – Electronic Structure of Atoms.
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Transcript of I. Waves & Particles Ch. 6 – Electronic Structure of Atoms.
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I. Waves & Particles
Ch. 6 – Electronic Structure of Atoms
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Properties of Waves
Many of the properties of light may be described in terms of waves even though light also has particle-like characteristics.
Waves are repetitive in nature
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A. Waves
Wavelength () - length of one complete wave; units of m or nm
Frequency () - # of waves that pass a point during a certain time period hertz (Hz) = 1/s
Amplitude (A) - distance from the origin to the trough or crest
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A. Waves
Agreater
amplitude
(intensity)
greater frequency
(color)
crest
origin
trough
A
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Electromagnetic Radiation
Electromagnetic radiation: (def) form of energy that exhibits wavelike behavior as it travels through space
Types of electromagnetic radiation: visible light, x-rays, ultraviolet (UV),
infrared (IR), radiowaves, microwaves, gamma rays
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Electromagnetic Spectrum
All forms of electromagnetic radiation move at a speed of about 3.0 x 108 m/s through a vacuum (speed of light)
Electromagnetic spectrum: made of all the forms of electromagnetic radiation
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B. EM Spectrum
LOW
ENERGY
HIGH
ENERGY
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B. EM Spectrum
LOW
ENERGY
HIGH
ENERGY
R O Y G. B I V
red orange yellow green blue indigo violet
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B. EM Spectrum
Frequency & wavelength are inversely proportional
c = c: speed of light (3.00 108 m/s): wavelength (m, nm, etc.): frequency (Hz)
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B. EM Spectrum
GIVEN:
= ?
= 434 nm = 4.34 10-7 m
c = 3.00 108 m/s
WORK: = c
= 3.00 108 m/s 4.34 10-7 m
= 6.91 1014 Hz
EX: Find the frequency of a photon with a wavelength of 434 nm.
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C. Quantum Theory
Photoelectric effect: emission of electrons from a metal when light shines on the metal
Hmm… (For a given metal, no electrons were emitted if the light’s frequency was below a certain minimum – why did light have to be of a minimum frequency?)
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C. Quantum Theory
Planck (1900)
Observed - emission of light from hot objects
Concluded - energy is emitted in small, specific amounts (quanta)
Quantum - minimum amount of energy change
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C. Quantum Theory
Planck (1900)
vs.
Classical Theory Quantum Theory
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C. Quantum Theory
Einstein (1905)
Observed - photoelectric effect
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C. Quantum Theory
E: energy (J, joules)h: Planck’s constant (6.626 10-34 J·s): frequency (Hz)
E = h
The energy of a photon is proportional to its frequency.
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C. Quantum Theory
GIVEN:
E = ? = 4.57 1014 Hzh = 6.6262 10-34 J·s
WORK:
E = h
E = (6.6262 10-34 J·s)(4.57 1014 Hz)
E = 3.03 10-19 J
EX: Find the energy of a red photon with a frequency of 4.57 1014 Hz.
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C. Quantum Theory
Einstein (1905)
Concluded - light has properties of both waves and particles
“wave-particle duality”
Photon - particle of light that carries a quantum of energy
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6.3. Bohr Model of the Atom
Ch.6-
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Excited and Ground State
Ground state: lowest energy state of an atom
Excited state: an atom has a higher potential energy than it had in its ground state
When an excited atom returns to its ground state, it gives off the energy it gained as EM radiation
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A. Line-Emission Spectrum
ground state
excited state
ENERGY IN PHOTON OUT
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B. Bohr Model
2) e- exist only in orbits with specific amounts of energy called energy levels
When e- are in these orbitals, they have fixed energy
Energy of e- are higher when they are further from the nucleus
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B. Bohr Model
Therefore…Bohr model leads us to conclude that:
e- can only gain or lose certain amounts of energy
only certain photons are produced
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B. Bohr Model
1
23
456 Energy of photon depends on the difference in energy levels
Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom
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C. Other Elementssummersummersummer
Each element has a unique bright-line emission spectrum.
“Atomic Fingerprint”
Helium
Bohr’s calculations only worked for hydrogen!
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III. Wave Behavior of Matter
Ch. 6 - Electrons in Atoms
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A. Electrons as Waves
Louis de Broglie (1924)
Applied wave-particle theory to e-
e- exhibit wave properties
QUANTIZED WAVELENGTHS
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A. Electrons as Waves
EVIDENCE: DIFFRACTION PATTERNS
ELECTRONSVISIBLE LIGHT
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A. Electrons as Waves
Diffraction: (def) bending of a wave as it
passes by the edge of an object
Interference: (def) when waves overlap (causes reduction and increase in energy in some areas of waves)
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6.5: Quantum Model
Chapter 6
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A. Quantum Mechanics
Heisenberg Uncertainty Principle
Impossible to know both the velocity and position of an electron
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A. Quantum Mechanics
σ3/2 Zπ
11s 0
eΨ a
Schrödinger Wave Equation (1926)
finite # of solutions quantized energy levels
defines probability of finding an e-
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B . Quantum Mechanics
Schrodinger wave equation and Heisenberg Uncertainty Principle laid foundation for modern quantum theory
Quantum theory: (def) describes mathematically the wave properties of e- and other very small particles
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B. Quantum Mechanics
Radial Distribution CurveOrbital
Orbital (“electron cloud”)
Region in space where there is 90% probability of finding an e-
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C. Quantum Numbers
UPPER LEVEL
Four Quantum Numbers:
Specify the “address” of each electron in an atom
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C. Quantum Numbers
1. Principal Quantum Number ( n )
Main energy level
Size of the orbital
n2 = # of orbitals in the energy level
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C. Quantum Numbers
s p d f
2. Angular Momentum Quantum # ( l ) Energy sublevel Shape of the orbital (# of possible shapes equal to n) values from 0 to n-1
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C. Quantum Numbers
If l equals… Then orbital shape is…
0 s
1 p
2 d
3 f
Principle quantum # followed by letter of sublevel
designates an atomic orbital
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C. Quantum Numbers
3. Magnetic Quantum Number ( ml )
Orientation of orbital
Specifies the exact orbitalwithin each sublevel
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C. Quantum Numbers
Values for ml:
m = -l… 0… +l
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C. Quantum Numbers
px py pz
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C. Quantum Numbers
Orbitals combine to form a spherical
shape.
2s
2pz2py
2px
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C. Quantum Numbers
4. Spin Quantum Number ( ms )
Electron spin +½ or -½
An orbital can hold 2 electrons that spin in opposite directions.
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C. Quantum Numbers
1. Principal # 2. Ang. Mom. # 3. Magnetic # 4. Spin #
energy level
sublevel (s,p,d,f)
orbital
electron
Pauli Exclusion Principle
No two electrons in an atom can have the same 4 quantum numbers.
Each e- has a unique “address”:
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C. Quantum Numbers
n = # of sublevels per level
n2 = # of orbitals per level
Sublevel sets: 1 s, 3 p, 5 d, 7 f
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Wrap-Up
Quantum # Symbol What it describes
Possible values
Principle quantum #
n main E level, size of orbital
n = positive whole integers
Angular Momentum Quantum #
l sublevels and their shapes
0 to (n-1)
Magnetic Quantum #
ml orientation of orbital
-l … 0 … +l
Spin Quantum #
ms
electron spin +1/2 or -1/2
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Electron Configuration
Ch. 6 - Electrons in Atoms
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a. ELECTRON CONFIGURATION
ELECTRON CONFIGURATION Notation to keep track of where electrons in an atom are distributed between shells and subshells
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B. General Rules
Pauli Exclusion Principle
Each orbital can hold TWO electrons
with opposite spins.
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B. General Rules
Aufbau Principle
Electrons fill the lowest energy orbitals first.
“Lazy Tenant Rule”
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RIGHTWRONG
B. General Rules
Hund’s Rule
Within a sublevel, place one e- per orbital before pairing them.
“Empty Bus Seat Rule”
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O
8e-
Orbital Diagram
Electron Configuration
1s2 2s2 2p4
C. Notation
1s 2s 2p
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Shorthand Configuration
S 16e-
Valence Electrons
Core Electrons
S 16e- [Ne] 3s2 3p4
1s2 2s2 2p6 3s2 3p4
C. Notation
Longhand Configuration
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© 1998 by Harcourt Brace & Company
sp
d (n-1)
f (n-2)
1234567
67
D. Periodic Patterns
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C. Periodic Patterns
Period # energy level (subtract for d & f)
A/B Group # total # of valence e-
Column within sublevel block # of e- in sublevel
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s-block
1st Period
1s11st column of s-block
C. Periodic Patterns
Example - Hydrogen
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1
2
3
4
5
6
7
C. Periodic Patterns
Shorthand Configuration Core e-: Go up one row and over to the
Noble Gas. Valence e-: On the next row, fill in the #
of e- in each sublevel.
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[Ar] 4s2 3d10 4p2
C. Periodic Patterns
Example - Germanium
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Full energy level
1
2
3
4 5
6
7
Full sublevel (s, p, d, f)Half-full sublevel
E. Stability
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Electron Configuration Exceptions
Copper
EXPECT: [Ar] 4s2 3d9
ACTUALLY: [Ar] 4s1 3d10
Copper gains stability with a full d-sublevel.
E. Stability
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Electron Configuration Exceptions
Chromium
EXPECT: [Ar] 4s2 3d4
ACTUALLY: [Ar] 4s1 3d5
Chromium gains stability with a half-full d-sublevel.
E. Stability
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E. Stability
Ion Formation Atoms gain or lose electrons to become
more stable. Isoelectronic with the Noble Gases.
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O2- 10e- [He] 2s2 2p6
E. Stability
Ion Electron Configuration
Write the e- config for the closest Noble Gas
EX: Oxygen ion O2- Ne