electrons in atoms. wavelength ( ) - length of one complete wave measured in m, cm, or nm in light...
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Electrons in Atoms
Wavelength () - length of one complete wave measured in m, cm, or nm
In light it tells us which color it is
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 (dip) or crest (peak Amount of energy the wave is carrying - height of wave. It
is measured in meters. In SOUND it tells us how LOUD it is. In LIGHT it tells how
BRIGHT it is.
A
greater amplitude
greater frequency
crest
origin
trough
A
Understanding electronic structure of atoms
Must understand light
Emitted and absorbed by substances. Visible light - type of Electromagnetic Radiation (EM)
Carries (radiant) energy through space Travels at speed of lightExhibits wavelike behavior.
Think of light as particlehelp understand how EM radiation and atomsinteract
LOW
ENERGY
HIGH
ENERGY
Move through a vacuum at the
‘speed of light’ 3.00 x 108 m/s
Behaves like waves that move through water
Result of energy transferred to the water (from
a stone)
Expressed as up and down movement
Both electric and magnetic properties
Wave Speed = (distance between wave peaks) x (frequency)
= (wavelength)
x (frequency)
EM radiation moves through a vacuum at the “speed of light”
3.00 x 108 m/s also called c. A lower energy wave (infrared and red) has a longer
wavelength() and lower frequency(f)
A higher energy wave (blue - violet) has a shorter
wavelength() and higher frequency(f).
Frequency & wavelength are inversely proportional
c = c: speed of light (3.00 108 m/s): wavelength (m, nm, etc.): frequency (Hz)
EX: Find the frequency of a photon with a wavelength of 434 nm.
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
Planck (1900)
Observed - emission of light from hot objects
Concluded - energy is emitted (absorbed or released) in small, specific amounts (quanta)
Quantum - smallest energy packet that can be emitted or absorbed as EM radiation by an atom.
E: energy (J, joules)h: Planck’s constant (6.6262 10-34 J·s): frequency (Hz)
E = h
Planck proposed that the energy, E, of a single quantum energy packet equals a constant (h) times its frequency
The energy of a photon is proportional to its frequency.
EX: Find the energy of a red photon with a frequency of 4.57 1014 Hz.
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
Energy is always emitted or absorbed in whole number multiples of hv, such as hv, 2 hv, 3 hv, 4hv, ….
The allowed energies are quantized values are restricted to certain quantities.
The notion of quantized rather than continuous energies is strange.
Ramp vs Staircase Ramp - vary the length your steps and energy used on
the walk up. Stairs - must exert exactly the specific amount of
energy needed to reach the next step. Your steps on steps are quantized, you cannot step
between them.
Planck (1900)
vs.
Classical Theory Quantum Theory
Einstein (1905)
Observed – photoelectric effect
Dispersed light falls on metal samples, the different frequencies produce different energetic photoelectrons
Einstein (1905)
Concluded - light has properties of both waves and particles (photons)
“wave-particle duality”
Photon - particle of light that carries a quantum of energy
Used planck’s quantum theory to deduced that: Ephoton = hv
Electrons in Atoms
ground state
excited state
ENERGY IN PHOTON OUT
Elements’ atoms absorb electrical energy
e- get excited, become unstable, and release energy
Energy is in form of light
Set of frequencies of EM waves
e- exist only in orbits with specific amounts of energy called energy levels
Therefore…
e- can only gain or lose certain amounts of energy
only certain photons are produced
Ground state = lowest allowable atomic electron energy state
Excited state = any higher energy state
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
1
23
456
Each element has a unique bright-line emission spectrum. “Atomic Fingerprint” Helium
Examples: Iron
Now, we can calculate for all elements and their electrons – next section
Louis de Broglie (1924)
Proposed eˉ in their orbits behave like a waveEVIDENCE: DIFFRACTION PATTERNS
ELECTRONSVISIBLE LIGHT
Heisenberg Uncertainty Principle Impossible to know velocity and position of an
electron at the same time
Trying to observe an electron’s position changes its momentum
Trying to observe an electron’s momentum changes its position
Electrons cannot be locked into well-defined circular orbits around the nucleus.
Orbital (“electron cloud”) a specific distribution of electron density in space. Each orbital has a characteristic energy and shape.
Orbital
Specify the “address” of each electron in an atom
UPPER LEVEL
1. Principal Quantum Number (n = 1, 2, 3, …) (see periodic table left column)
Indicates the relative size and energy of atomic orbitals
As (n) increases, the orbital becomes larger, the electron spends more time farther from the nucleus
Each major energy level is called a principle energy level Ex: lowest level = 1 ground state, highest level = 7 excited state
2. Energy Sublevel Defines the shape of the orbital (s, p, d, f)
# of orbital related to each sublevel is always an odd # s = 1, p = 3, d = 5, f = 7
Each orbital can contain at most 2 electrons
s p d f
Subscripts x, y, z designates orientation Specifies the exact orbital within each
sublevel
px py pz
4. Spin Quantum Number ( ms ) Electron spin +½ or -½
An orbital can hold 2 electrons that spin in opposite directions.
Pauli Exclusion Principle A maximum of 2 electrons can occupy a single
atomic orbital
Only if they have opposite spins
1. Principal #
2. Energy sublevel 3. Orientation
4. Spin #
energy level
(s,p,d,f)
x, y, z
exact electron
Electron Configuration
Electrons in Atoms
A. General Rules
Aufbau Principle
Electrons fill the lowest energy orbitals first.
“Lazy Tenant Rule”
RIGHTWRONG
A. 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
B. Notation
1s 2s 2p
Shorthand Configuration
S 16e-
Valence ElectronsCore Electrons
S 16e- [Ne] 3s2 3p4
1s2 2s2 2p6 3s2 3p4
B. Notation
Longhand Configuration
Valence electrons: determine chemical properties of that
element & are the electrons in the atoms outermost orbital
© 1998 by Harcourt Brace & Company
sp
d (n-1)
f(n-2)
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Notation
1
2
3
4
5
6
7
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.
Shorthand Notation
[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
D. Stability
Electron Configuration Exceptions
Copper
EXPECT: [Ar] 4s2 3d9
ACTUALLY: [Ar] 4s1 3d10
Copper gains stability with a full d-sublevel.
D. Stability
Electron Configuration Exceptions
Chromium
EXPECT: [Ar] 4s2 3d4
ACTUALLY: [Ar] 4s1 3d5
Chromium gains stability with a half-full d-sublevel.
D. Stability
D. Stability
Ion Formation Atoms gain or lose electrons to become
more stable. Isoelectronic with the Noble Gases.
O2- 10e- [He] 2s2 2p6
D. Stability
Ion Electron Configuration
Write the e- config for the closest Noble Gas
EX: Oxygen ion O2- Ne