i. waves & particles ch. 6 – electronic structure of atoms
TRANSCRIPT
I. Waves & Particles
Ch. 6 – Electronic Structure of Atoms
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
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
A. Waves
Agreater
amplitude
(intensity)
greater frequency
(color)
crest
origin
trough
A
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
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
B. EM Spectrum
LOW
ENERGY
HIGH
ENERGY
B. EM Spectrum
LOW
ENERGY
HIGH
ENERGY
R O Y G. B I V
red orange yellow green blue indigo violet
B. EM Spectrum
Frequency & wavelength are inversely proportional
c = c: speed of light (3.00 108 m/s): wavelength (m, nm, etc.): frequency (Hz)
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.
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?)
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
C. Quantum Theory
Planck (1900)
vs.
Classical Theory Quantum Theory
C. Quantum Theory
Einstein (1905)
Observed - photoelectric effect
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.
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.
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
6.3. Bohr Model of the Atom
Ch.6-
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
A. Line-Emission Spectrum
ground state
excited state
ENERGY IN PHOTON OUT
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
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
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
C. Other Elementssummersummersummer
Each element has a unique bright-line emission spectrum.
“Atomic Fingerprint”
Helium
Bohr’s calculations only worked for hydrogen!
III. Wave Behavior of Matter
Ch. 6 - Electrons in Atoms
A. Electrons as Waves
Louis de Broglie (1924)
Applied wave-particle theory to e-
e- exhibit wave properties
QUANTIZED WAVELENGTHS
A. Electrons as Waves
EVIDENCE: DIFFRACTION PATTERNS
ELECTRONSVISIBLE LIGHT
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)
6.5: Quantum Model
Chapter 6
A. Quantum Mechanics
Heisenberg Uncertainty Principle
Impossible to know both the velocity and position of an electron
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-
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
B. Quantum Mechanics
Radial Distribution CurveOrbital
Orbital (“electron cloud”)
Region in space where there is 90% probability of finding an e-
C. Quantum Numbers
UPPER LEVEL
Four Quantum Numbers:
Specify the “address” of each electron in an atom
C. Quantum Numbers
1. Principal Quantum Number ( n )
Main energy level
Size of the orbital
n2 = # of orbitals in the energy level
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
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
C. Quantum Numbers
3. Magnetic Quantum Number ( ml )
Orientation of orbital
Specifies the exact orbitalwithin each sublevel
C. Quantum Numbers
Values for ml:
m = -l… 0… +l
C. Quantum Numbers
px py pz
C. Quantum Numbers
Orbitals combine to form a spherical
shape.
2s
2pz2py
2px
C. Quantum Numbers
4. Spin Quantum Number ( ms )
Electron spin +½ or -½
An orbital can hold 2 electrons that spin in opposite directions.
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”:
C. Quantum Numbers
n = # of sublevels per level
n2 = # of orbitals per level
Sublevel sets: 1 s, 3 p, 5 d, 7 f
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
Electron Configuration
Ch. 6 - Electrons in Atoms
a. ELECTRON CONFIGURATION
ELECTRON CONFIGURATION Notation to keep track of where electrons in an atom are distributed between shells and subshells
B. General Rules
Pauli Exclusion Principle
Each orbital can hold TWO electrons
with opposite spins.
B. General Rules
Aufbau Principle
Electrons fill the lowest energy orbitals first.
“Lazy Tenant Rule”
RIGHTWRONG
B. General Rules
Hund’s Rule
Within a sublevel, place one e- per orbital before pairing them.
“Empty Bus Seat Rule”
O
8e-
Orbital Diagram
Electron Configuration
1s2 2s2 2p4
C. Notation
1s 2s 2p
Shorthand Configuration
S 16e-
Valence Electrons
Core Electrons
S 16e- [Ne] 3s2 3p4
1s2 2s2 2p6 3s2 3p4
C. Notation
Longhand Configuration
© 1998 by Harcourt Brace & Company
sp
d (n-1)
f (n-2)
1234567
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D. Periodic Patterns
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
s-block
1st Period
1s11st column of s-block
C. Periodic Patterns
Example - Hydrogen
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.
[Ar] 4s2 3d10 4p2
C. Periodic Patterns
Example - Germanium
Full energy level
1
2
3
4 5
6
7
Full sublevel (s, p, d, f)Half-full sublevel
E. Stability
Electron Configuration Exceptions
Copper
EXPECT: [Ar] 4s2 3d9
ACTUALLY: [Ar] 4s1 3d10
Copper gains stability with a full d-sublevel.
E. Stability
Electron Configuration Exceptions
Chromium
EXPECT: [Ar] 4s2 3d4
ACTUALLY: [Ar] 4s1 3d5
Chromium gains stability with a half-full d-sublevel.
E. Stability
E. Stability
Ion Formation Atoms gain or lose electrons to become
more stable. Isoelectronic with the Noble Gases.
O2- 10e- [He] 2s2 2p6
E. Stability
Ion Electron Configuration
Write the e- config for the closest Noble Gas
EX: Oxygen ion O2- Ne