Transcript Matter, Measurements and Problem Solving

Chapter 6
Electronic Structure of Atoms
SC 131 CHEM 1
Chemistry: The Central Science
CM Lamberty
Quantum Mechanics: A Theory
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Smallness of atoms and subatomic particles
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Traditional observations not possible
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Size of e- <10-9 of 10-9 of a gram
Speck of dust contains as many e- as there have been
people on Earth since beginning
Dtm chemical and physical properties
e- do not move in regular patterns
e- observed behave differently than those not observed.
Quantum-mechanical model
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Model to explain how e- exist in atoms and dtm
properties
Explain WHY some M, some NM, why noble gases are
inert, etc.
The Nature of Light
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Wave Nature of Light
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The Electromagnetic Spectrum
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Radio (low E) to Gamma rays (high E)
Interference and Diffraction
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Properties of waves
Ways waves may interact
The Particle Nature of Light
Photoelectric effect
 photons
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The Wave Nature of Matter
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Light is electromagnetic
radiation
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Wave composed of
oscillating mutually
perpendicular electric and
magnetic fields
Speed of light (vacuum)
3.00x108 m/s
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Amplitude
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Vertical height of crest
Determines the
intensity of light
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The Wave Nature of Matter
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Wavelength
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Distance between adjacent crests
Frequency
Number of cycles passing a point in given
period of time
 Cycles per second (s-1). 1 Hertz = 1 cycle/s
 Frequency directly proportional to speed,
inverse to wavelength
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n= c
l
Wavelength and Amplitude
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The Electromagnetic Spectrum
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Includes ALL wavelengths of EM radiation
10-15m (gamma) - 105m (radio waves)
Short wavelength has greater E
Gamma (g) rays
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Most energetic, shortest
Produced by sun and stars and unstable atomic nuclei
Damage to biological molecules
X-rays
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Longer wavelength than gamma
Pass through many substances that block visible
Can damage biological molecules
The Electromagnetic Spectrum
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Ultraviolet
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Visible
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Component of sunlight for suntan/sunburn
Carries enough E to damage biological mq
excessive exposure skin cancer, cataracts
Violet (short l, high E) - red (longer l, lower E)
Violet, blue, green, yellow, orange, red
Causes certain mq in eye to change shape resulting in
vision
Color we see is reflected, others absorbed
Infrared
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Heat from hot object
Night vision goggles
The Electromagnetic Spectrum
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Microwaves
Longer wavelengths
 Used for radar and microwave ovens
 Efficiently absorbed by water and can heat
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Radio waves
Longest wavelength
 Transmit signals responsible for FM and AM
radio, cellular phones, TV, etc
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Interference and Diffraction
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Interference
How waves add together
 Constructive or destructive
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Diffraction
How waves bend to move around/though
object
 Diffraction of light through 2 slits
 Interference pattern
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Interference and Diffraction
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The Particle Nature of Light
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Light initially thought of as wave
Photoelectric effect
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Metals emit e- when light shines on them
Series of tests did not follow EM theory
Einstein: packets of light E = hn
h is Planck’s constant
Photons
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Our name for packets of light
Sometimes called quantum of light
E = hc
l
Light is “lumpy”
Light is shower of particles each having e of hn
Wave-particle duality of light
Atomic Spectroscopy & Bohr Model
Study of the EM radiation absorbed and
emitted by atoms
 Atom absorbs E (heat, light, electricity)
and remits the E as light
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Each element emit light of characteristic color
 Each with several distinct wavelengths
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Emission spectrum
Each element has its own emission spectrum
 Discrete lines not continuous
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Atomic Spectroscopy & Bohr Model
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Atomic Spectroscopy & Bohr Model
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Johannes Rydberg
Simple equation to predict wavelength of H
 1/l = R(1/m2-1/n2)
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Neils Bohr
His model: e- travel around nucleus in circular
orbits.
 These orbits can exist only as specific fixed
distances from nucleus
 E of each orbit was fixed or quantized
 Stationary states
 Only when e- made a transition that radiation
emitted or absorbed
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The Wave Nature of Matter
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Louis de Broglie
Wave nature of electrons
 Diffraction pattern
 de Broglie relation l = h/mn
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Heisenberg
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Uncertainty Principle: cannot simultaneously
observe both the wave nature and the particle
nature of the electron
Quantum Mechanics and the Atom
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Schrodinger
Orbital, probability distribution map showing
where the electron is likely to be found
 Wave function
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Quantum Numbers used to specify each
orbital or location of electron for an atom.
Quantum Mechanics and the Atom
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Principle quantum number, n
Integer that dtm overall size and E of orbital
 n= 1,2,3…
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Angular quantum number, l
Integer that dtm shape of orbital
 l = 0,1,2,…(n-1)
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Magnetic quantum number, ml
Integer that dtm orientation of orbital
 ml = -l to +l (-l, …, -1, 0, 1,…, l)
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Quantum Mechanics and the Atom
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Atomic Spectroscopy Explained
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The Shapes of Atomic Orbitals
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Shape important b/c covalent chemical
bonds depend upon sharing of electrons
and occupy these orbitals
Shapes of the overlapping orbitals dtm shape
of molecule
 Shape dtm primarily by l the angular
momentum quantum number
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l=0
 l=1
 l=2
 l=3
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s orbital
p orbital
d orbital
f orbital
The Shapes of Atomic Orbitals
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s orbitals
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The Shapes of Atomic Orbitals
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p orbitals
2 lobes
 Node at nucleus
 Orbitals are orthogonal to one another
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The Shapes of Atomic Orbitals
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d orbitals
5 3d orbitals
 4 are cloverleaf with 4 lobes
 5th is 2-lobed with donut (see p. 266)
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Electron Configurations
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Electron configuration
Ground state
Electron spin and Pauli Exclusion Principle
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Direction of arrow represents electron spin
Direction does not affect value
Direction is quantized either up or down
Spin Quantum Number, ms
+1/2 (up) or -1/2 (down)
Electron Configurations
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Pauli Exclusion Principle
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No two electrons can have the same four quantum
numbers
Electron Configurations
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Sublevel Energy Splitting in Multielectron Atoms
E(s) < E(p) < E(d) < E(f)
Sheilding
Effective nuclear charge
Electron Configuration of Sulfur