Chapter 5 Electrons In Atoms. Topics to Be Covered 5.1 Light and Quantized Energy 136-145 5.2...
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Transcript of Chapter 5 Electrons In Atoms. Topics to Be Covered 5.1 Light and Quantized Energy 136-145 5.2...
![Page 1: Chapter 5 Electrons In Atoms. Topics to Be Covered 5.1 Light and Quantized Energy 136-145 5.2 Quantum Theory and the Atom 146-155 5.3 Electron Configuration.](https://reader036.fdocuments.us/reader036/viewer/2022062308/56649cec5503460f949b8a08/html5/thumbnails/1.jpg)
Chapter 5
Electrons In Atoms
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Topics to Be Covered
5.1 Light and Quantized Energy 136-145
5.2 Quantum Theory and the Atom 146-155
5.3 Electron Configuration 156-162
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Section 5.1
Light and Quantized Energy
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The Atom & Unanswered Questions
Early 1900s Discovered 3 subatomic particles Continued quest to understand atomic
structure Rutherford’s model
Positive charge in nucleus Fast moving electrons around that No accounting for differences and
similarities in chemical behavior
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The Atom and Unanswered Questions
Example: Lithium, sodium, and potassium have
similar chemical behaviors (explained more in next chapter)
Early 1900s Scientists began to unravel mystery Certain elements emitted visible light
when heated in a flame Analysis revealed chemical behavior
depends on arrangement of electrons
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The Wave Nature of Light
Electromagnetic radiation A form of energy that exhibits wavelike
behavior as it travels through space Visible light is a type of ER
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Characteristics of Waves
All waves can be described by several characteristics
1. Wavelength2. Frequency3. Amplitude
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Wavelength
Represented by lambda λ Shortest distance between
equivalent points on a continuous waves
Measure crest to crest or trough to trough
Usually expressed in m, cm, or nm
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Frequency
Represented by nu ν The number of waves that pass a
given point per second Given in the unit of hertz (Hz) 1 Hz = 1 wave per second
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Amplitude
The wave’s height from the origin to a crest or from the origin to a trough
Wavelength and frequency do not affect amplitude
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Speed
All electromagnetic waves in a vacuum travel at a speed of 3.00 x 108 m/s This includes visible light
The speed of light has its own symbol C
C= λν
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Electromagnetic Spectrum
Also called the EM spectrum Includes all forms of
electromagnetic radiation With the only differences in the
types of radiation being their frequencies and wavelengths
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Electromagnetic Spectrum
Figure 5.5
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Problems
Page 140 Calculating Wavelength of an EM
Wave
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Particle Nature of Light
Needed to explain other properties of light Heated objects emit only certain
frequencies of light at a given temperature
Some metals emit electrons when light of a specific frequency shines on them
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Quantum Concept
When objects are heated they emit glowing light
1900 Max Planck began searching for an
explanation Studied the light emitted by heated
objects Startling conclusion
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Quantum Concept
Planck discovered: Matter can gain or lose energy only in
small specific amounts These amounts are called quanta Quantum—is the minimum amount of
energy that can be gained or lost by an atom
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Example
Heating a cup of water Most people thought that you can add
any amount of thermal energy to the water by regulating the power and duration of the microwaves
In actuality, the temperature increases in infinitesimal steps as its molecules absorb quanta of energy, which appear to be a continuous manner
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Quantum Concept
Planck proposed that energy emitted by hot objects was quantized
Planck further demonstrated mathematically that a relationship exists between energy of a quantum and a frequency
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Energy of a Quantum
Equantum=hv
Equantum represents energy h is Planck’s constant v represents frequency
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Planck’s Constant
Symbol = h 6.626 x 10-34 J*s J is the symbol for joule
The SI unit of energy The equation shows that the energy
of radiation increases as the radiation’s frequency, v, increases.
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Planck’s Theory
For given frequencies Matter can emit/absorb energy only in
whole number multiples of hv 1hv, 2hv, 3hv, 4hv etc.
Matter can have only certain amounts of energy
Quantities of energy between these values do not exist
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The Photoelectric Effect
Photoelectric effect electrons, called photoelectrons are emitted from a metal’s surface when light of a certain frequency, or
higher than a certain frequency shines on the surface
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Light’s Dual Nature
Einstein proposed in 1905 that light has a dual nature
photon—a massless particle that carries a quantum of energy
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Energy of a Photon
Ephoton=hv
Ephoton represents energy h is Planck’s constant v represents frequency
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Light’s Dual Nature
Einstein proposed Energy of a photon must have a certain
threshold value to cause the ejection of a photoelectron from the surface of a metal
Even small #s of photons with energy above the threshold value will cause the photoelectric effect
Einstein won Nobel Prize in Physics in 1921
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Sample Problems
Page 143 Sample Problem 5.2 Calculating Energy of a Photon
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Atomic Emission Spectra
See page 145
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Section 5.2
Quantum Theory and The Atom
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Bohr’s Model of the Atom
Dual-nature explains more Atomic Emission Spectra
Not continuous Only certain frequencies of light
Explained the Atomic Emission Spectra
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Energy States of Hydrogen
Bohr proposed certain allowable energy states
Bohr proposed electrons could travel in certain orbitals
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Energy states of Hydrogen
Ground State Lowest allowable energy state of an
atom Orbit size
Smaller the orbit, the lower the energy state/level
Larger the orbit, the higher the energy state/level
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Energy states of Hydrogen
Hydrogen can have many excited states It only has one electron
Quantum Number Number assigned to each orbital n Look at Table 5.1
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The Hydrogen Line Spectrum
Hydrogen Ground State Electron is in n=1 orbit Does not radiate energy
Hydrogen Excited State Energy is added to the atom from
outside source Electron moves to a higher energy orbit
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The Hydrogen Line Spectrum
Only Certain Atomic Energy Levels Possible
Example Our Classroom
Balmer Series Electron transitions from higher-energy
orbits to the second orbit Account for visible lines
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The Hydrogen Line Spectrum
Lyman Series Ultraviolet Electrons drop into n=1 orbit
Paschen Series Infrared Electrons drop into n = 3 orbit