how did we discover electron arrangement in an atom? electromagnetic radiation ! ! !
TRANSCRIPT
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How did we discover electron arrangement
in an atom?
ELECTROMAGNETIC
RADIATION ! ! !
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Waves
Repeated disturbance through a medium (air, liquid) from origin to distant points.
Medium does not move
Ex. Ocean waves, sound waves
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Characteristics of Waves
Wavelength Distance between 2 points within a wave cycle 2 peaks
Frequency # of wave cycles passing a point for a
particular time unit Usually seconds.
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Wavelength and frequency are inversely proportional.
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c = νλ
c = speed of light, 3.0 x 10 8 m/s Constant
ν= frequency (s-1 or Hz)
λ= wavelength (m)
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Example 1:
Find the frequency of a green light that has a wavelength of 545 nm.
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Electromagnetic Waves
Produced from electric charge movement
Changes within electric and magnetic fields carried over a distance
No medium needed
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Electromagnetic Spectrum
Contains full range of wavelengths and frequencies found with electromagnetic radiation Wavelength/frequency changes cause color
changes
Mostly invisible, visible range (390 nm -760 nm)
Different materials absorb/transmit the spectrum differently.
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Types of Spectra What is a spectra?
Spectrum– white light/radiation split into different wavelengths and frequencies by a prism
Continuous spectrum No breaks in spectrum Colors together
Line spectrum Line pattern emitted by light from excited atoms of
a particular element Aided in determining atomic structure
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Line Spectrum
Pattern emitted by light from excited atoms of an element
Specific for each element Only certain wavelengths of visible spectrum
present
Used for element identification
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Flame Tests
Some atoms of elements produce visible light if heated
Each element has a specific flame color
Examples: Li, Na, Cs, Ca
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A Bit of Quantum Theory……
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Max Planck 1900
Related energy and radiation E = hν h= 6.626 x 10 -34 Js (Planck’s constant) E = energy per photon (J)
Quantum---smallest amount of energy
Atoms can only absorb/emit specific quanta
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Albert Einstein 1905
Added to Planck’s concept
Photons— Bundles of light energy Same energy as quantum E = hν (energy of photon)
Photons release energy and electrons gain energy Threshold frequency– minimum amount of energy
needed by photon to extract electron
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THEREFORE ………
Light is in the form of electromagnetic waves
Photons can resemble particles
Gave raise to the possibility of thinking about wave AND particle qualities of subatomic particles (electron)
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Example 1
Calculate the energy found in a photon of red light with a wavelength of 700.0 nm
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Example 2
How much energy (in joules) is found in the radiation of the hydrogen atom emission spectrum with a 656.3 nm wavelength?
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Example 3:
A sodium atom emits yellow light with a wavelength of 589 nm when it is excited. Find the energy per photon of this light.
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Quantum Theory/Electron
Orbitals
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Coulomb’s Law
Describes the attractive force between negative electrons and positive nucleus. Force is directly related to the charge of electron and
nucleus Force is inversely related to distance between particles
F = qe x qp
r2
(IE…an electron’s energy is dependent on distance from nucleus)
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Early Models of the Atom Bohr
1913—hydrogen atom structure
Physics + quantum theory
Electrons move in definite orbits around the positively charged nucleus—planetary model
Does not apply as atoms increase in electron number
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Electrons orbit nucleus in different energy levels
Lower energy levels, closest to nucleus (n = 1)
Higher energy levels increase electron’s distance from nucleus
Electrons can “transition” or jump between energy levels through photons Gain/absorb photon—higher energy level Lose/emit photon—lower energy level
Bohr Model
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Energy States in an Atom
Atoms can gain or loss energy.
Specific energy states within an atom. Can be counted Ground State = lowest energy state Excited State = higher energy level
than ground, gained energy
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So, where does the Bohr Model fit in?
Electrons orbit around the nucleus at different energy levels/orbits.
Electron’s energy level = orbit level where electron is located.
Light absorption = electron moves from a state of low energy to high energy. “becomes excited”
Light Emitted = electron falls from an “excited” state of energy to a lower energy level.
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Ex. Li
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Erwin Schrödinger
Quantum mechanics
1926---wave equation
Electrons behave more like waves than particles
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Heisenberg’s Uncertainty Principle Electron’s location and direction cannot be
known simultaneously
Electron as cloud of negative charge
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Modern Model of the AtomThe electron cloud
Sometimes called the wave model
Electron as cloud of negative charge
Spherical cloud of varying density
Varying density shows where an electron is more or less likely to be
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Treats electron’s location as wave property
Defined by quantum numbers
Quantum numbers Provide information about size, shape, and
orientation of atomic orbitals Define atomic orbitals from general to specific
Quantum Theory
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Determines orbital size and electron energy
Same as “n” value/orbital in Bohr model
Positive whole number, NOT 0
Shells – orbitals with same value
n = 1, 2, 3, 4, etc.
Principal Quantum Number (n)
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Defines orbital shape for a particular region of atom
Think as “subshell”
l = n-1
# of orbitals/subshells = principal quantum #
Orbital Angular Momentum Quantum
Number (l)
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l Orbital/Subshell
0 s
1 p
2 d
3 f
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Describes orbital orientation within an atom
Range from –l to +l, 0 is possible
ml = 0, ± 1, ± 2, etc.
ml = 2l + 1 (number of orientations)
Magnetic Quantum Number (ml)
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2p
4f
How do you specify orbitals?
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s orbital 1 possible orbital orientation, spherical shape n value determines size Charge cloud found near center, likely electron location
p orbital 3 possible orbital orientations, dumbbell shape pX, py, pz
Orbital Shapes
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P
What does atomic structure REALLY look
like?
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Describes the motion of an electron, spinning
As electron moves, magnetic field induced
Electrons with opposite spins, cancel magnetic field of other
Values: +1/2, -1/2
Electron Spin
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Read lab procedure
Read pp. 267-289 (for Friday)
Homework