I. Waves & Particles Ch. 4 - Electrons in Atoms. A. Waves zWavelength ( ) - length of one complete...
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Transcript of I. Waves & Particles Ch. 4 - Electrons in Atoms. A. Waves zWavelength ( ) - length of one complete...
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I. Waves & Particles
Ch. 4 - Electrons in Atoms
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A. Waves
Wavelength () - length of one complete wave
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
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A. Waves
Agreater
amplitude
(intensity)
greater frequency
(color)
crest
origin
trough
A
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B. EM Spectrum
LOW
ENERGY
HIGH
ENERGY
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B. EM Spectrum
LOW
ENERGY
HIGH
ENERGY
R O Y G. B I V
red orange yellow green blue indigo violet
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B. EM Spectrum
Frequency & wavelength are inversely proportional
c = c: speed of light (3.00 108 m/s): wavelength (m, nm, etc.): frequency (Hz)
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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.
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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
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C. Quantum Theory
Planck (1900)
vs.
Classical Theory Quantum Theory
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C. Quantum Theory
Einstein (1905)
Observed - photoelectric effect
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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
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C. Quantum Theory
E: energy (J, joules)h: Planck’s constant (6.6262 10-34 J·s): frequency (Hz)
E = h
The energy of a photon is proportional to its frequency.
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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.
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II. Bohr Model of the Atom
Ch. 4 - Electrons in Atoms
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A. Line-Emission Spectrum
ground state
excited state
ENERGY IN PHOTON OUT
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B. Bohr Model
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
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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
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C. Other Elements
Each element has a unique bright-line emission spectrum.
“Atomic Fingerprint”
Helium
Bohr’s calculations only worked for hydrogen!
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III. Quantum Model
of the Atom
Ch. 4 - Electrons in Atoms
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A. Electrons as Waves
Louis de Broglie (1924)
Applied wave-particle theory to e-
e- exhibit wave properties
QUANTIZED WAVELENGTHS
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A. Electrons as Waves
QUANTIZED WAVELENGTHS
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A. Electrons as Waves
EVIDENCE: DIFFRACTION PATTERNS
ELECTRONSVISIBLE LIGHT
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B. Quantum Mechanics
Heisenberg Uncertainty Principle
Impossible to know both the velocity and position of an electron at the same time
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B. 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-
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B. Quantum Mechanics
Radial Distribution CurveOrbital
Orbital (“electron cloud”)
Region in space where there is 90% probability of finding an e-
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C. Quantum Numbers
UPPER LEVEL
Four Quantum Numbers:
Specify the “address” of each electron in an atom
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C. Quantum Numbers
1. Principal Quantum Number ( n )
Energy level
Size of the orbital
n2 = # of orbitals in the energy level
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C. Quantum Numbers
s p d f
2. Angular Momentum Quantum # ( l )
Energy sublevel
Shape of the orbital
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C. Quantum Numbers
n = # of sublevels per level
n2 = # of orbitals per level
Sublevel sets: 1 s, 3 p, 5 d, 7 f
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C. Quantum Numbers
3. Magnetic Quantum Number ( ml )
Orientation of orbital
Specifies the exact orbitalwithin each sublevel
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C. Quantum Numbers
px py pz
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C. Quantum Numbers
Orbitals combine to form a spherical
shape.
2s
2pz2py
2px
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C. Quantum Numbers
4. Spin Quantum Number ( ms )
Electron spin +½ or -½
An orbital can hold 2 electrons that spin in opposite directions.
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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”:
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Feeling overwhelmed?
Read Section 4-2!
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IV. Electron Configuration(p. 105 - 116,
128 - 139)
Ch. 4 - Electrons in Atoms
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A. General Rules
Pauli Exclusion Principle
Each orbital can hold TWO electrons
with opposite spins.
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A. General Rules
Aufbau Principle
Electrons fill the lowest energy orbitals first.
“Lazy Tenant Rule”
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RIGHTWRONG
A. General Rules
Hund’s Rule
Within a sublevel, place one e- per orbital before pairing them.
“Empty Bus Seat Rule”
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O
8e-
Orbital Diagram
Electron Configuration
1s2 2s2 2p4
B. Notation
1s 2s 2p
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Shorthand Configuration
S 16e-
Valence Electrons
Core Electrons
S 16e- [Ne] 3s2 3p4
1s2 2s2 2p6 3s2 3p4
B. Notation
Longhand Configuration
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© 1998 by Harcourt Brace & Company
sp
d (n-1)
f (n-2)
1234567
67
C. Periodic Patterns
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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
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s-block
1st Period
1s11st column of s-block
C. Periodic Patterns
Example - Hydrogen
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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.
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[Ar] 4s2 3d10 4p2
C. Periodic Patterns
Example - Germanium
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Full energy level
1
2
3
4 5
6
7
Full sublevel (s, p, d, f)Half-full sublevel
D. Stability
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Electron Configuration Exceptions
Copper
EXPECT: [Ar] 4s2 3d9
ACTUALLY: [Ar] 4s1 3d10
Copper gains stability with a full d-sublevel.
D. Stability
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Electron Configuration Exceptions
Chromium
EXPECT: [Ar] 4s2 3d4
ACTUALLY: [Ar] 4s1 3d5
Chromium gains stability with a half-full d-sublevel.
D. Stability
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D. Stability
Ion Formation Atoms gain or lose electrons to become
more stable. Isoelectronic with the Noble Gases.
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O2- 10e- [He] 2s2 2p6
D. Stability
Ion Electron Configuration
Write the e- config for the closest Noble Gas
EX: Oxygen ion O2- Ne
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Ch. 5 - The Periodic Table
I. History
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A. Mendeleev
Dmitri Mendeleev (1869, Russian) Organized elements
by increasing atomic mass.
Elements with similar properties were grouped together.
There were some discrepancies.
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A. Mendeleev
Dmitri Mendeleev (1869, Russian) Predicted properties of undiscovered
elements.
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B. Moseley
Henry Mosely (1913, British)
Organized elements by increasing atomic number.
Resolved discrepancies in Mendeleev’s arrangement.
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II. Organization of theElements
Ch. 5 - The Periodic Table
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MetalsNonmetalsMetalloids
A. Metallic Character
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Main Group ElementsTransition MetalsInner Transition Metals
B. Blocks
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III. Periodic Trends
Ch. 5 - The Periodic Table
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0 5 10 15 20Atomic Number
Ato
mic
Ra
diu
s (
pm
)
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A. Periodic Law
When elements are arranged in order of
increasing atomic #, elements with similar
properties appear at regular intervals.
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0 5 10 15 20Atomic Number
Ato
mic
Ra
diu
s (
pm
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B. Chemical Reactivity
Families Similar valence e- within a group result in
similar chemical properties
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B. Chemical Reactivity
Alkali MetalsAlkaline Earth MetalsTransition MetalsHalogensNoble Gases
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Atomic Radius size of atom
© 1998 LOGALFirst Ionization Energy
Energy required to remove one e- from a neutral atom.
© 1998 LOGAL
Melting/Boiling Point
C. Other Properties
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Atomic Radius
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Ato
mic
Ra
diu
s (
pm
)
D. Atomic Radius
Li
ArNe
KNa
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Atomic Radius Increases to the LEFT and DOWN
D. Atomic Radius
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Why larger going down?
Higher energy levels have larger orbitals
Shielding - core e- block the attraction between the nucleus and the valence e-
Why smaller to the right?
Increased nuclear charge without additional shielding pulls e- in tighter
D. Atomic Radius
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First Ionization Energy
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t Io
niz
ati
on
En
erg
y (k
J)
E. Ionization Energy
KNaLi
Ar
NeHe
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First Ionization Energy Increases UP and to the RIGHT
E. Ionization Energy
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Why opposite of atomic radius?
In small atoms, e- are close to the nucleus where the attraction is stronger
Why small jumps within each group?
Stable e- configurations don’t want to lose e-
E. Ionization Energy
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Successive Ionization Energies
Mg 1st I.E. 736 kJ
2nd I.E. 1,445 kJ
Core e- 3rd I.E. 7,730 kJ
Large jump in I.E. occurs when a CORE e- is removed.
E. Ionization Energy
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Al 1st I.E. 577 kJ
2nd I.E. 1,815 kJ
3rd I.E. 2,740 kJ
Core e- 4th I.E. 11,600 kJ
Successive Ionization Energies
Large jump in I.E. occurs when a CORE e- is removed.
E. Ionization Energy
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Melting/Boiling Point Highest in the middle of a period.
F. Melting/Boiling Point
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Ionic Radius
Cations (+)
lose e-
smaller
© 2002 Prentice-Hall, Inc.
Anions (–)
gain e-
larger
G. Ionic Radius
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Which atom has the larger radius?
Be or Ba
Ca or Br
Ba
Ca
Examples
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Which atom has the higher 1st I.E.?
N or Bi
Ba or Ne
N
Ne
Examples
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Which atom has the higher melting/boiling point?
Li or C
Cr or Kr
C
Cr
Examples
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Which particle has the larger radius?
S or S2-
Al or Al3+
S2-
Al
Examples
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Periodic Trends Summary