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I. Waves & Particles
Ch. 4 - Electrons in Atoms
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
A. Waves
Agreater
amplitude
(intensity)
greater frequency
(color)
crest
origin
trough
A
B. EM Spectrum
LOW
ENERGY
HIGH
ENERGY
B. EM Spectrum
LOW
ENERGY
HIGH
ENERGY
R O Y G. B I V
red orange yellow green blue indigo violet
B. EM Spectrum
Frequency & wavelength are inversely proportional
c = c: speed of light (3.00 108 m/s): wavelength (m, nm, etc.): frequency (Hz)
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.
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
C. Quantum Theory
Planck (1900)
vs.
Classical Theory Quantum Theory
C. Quantum Theory
Einstein (1905)
Observed - photoelectric effect
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
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.
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.
II. Bohr Model of the Atom
Ch. 4 - Electrons in Atoms
A. Line-Emission Spectrum
ground state
excited state
ENERGY IN PHOTON OUT
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
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
C. Other Elements
Each element has a unique bright-line emission spectrum.
“Atomic Fingerprint”
Helium
Bohr’s calculations only worked for hydrogen!
III. Quantum Model
of the Atom
Ch. 4 - Electrons in Atoms
A. Electrons as Waves
Louis de Broglie (1924)
Applied wave-particle theory to e-
e- exhibit wave properties
QUANTIZED WAVELENGTHS
A. Electrons as Waves
QUANTIZED WAVELENGTHS
A. Electrons as Waves
EVIDENCE: DIFFRACTION PATTERNS
ELECTRONSVISIBLE LIGHT
B. Quantum Mechanics
Heisenberg Uncertainty Principle
Impossible to know both the velocity and position of an electron at the same time
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-
B. Quantum Mechanics
Radial Distribution CurveOrbital
Orbital (“electron cloud”)
Region in space where there is 90% probability of finding an e-
C. Quantum Numbers
UPPER LEVEL
Four Quantum Numbers:
Specify the “address” of each electron in an atom
C. Quantum Numbers
1. Principal Quantum Number ( n )
Energy level
Size of the orbital
n2 = # of orbitals in the energy level
C. Quantum Numbers
s p d f
2. Angular Momentum Quantum # ( l )
Energy sublevel
Shape of the orbital
C. Quantum Numbers
n = # of sublevels per level
n2 = # of orbitals per level
Sublevel sets: 1 s, 3 p, 5 d, 7 f
C. Quantum Numbers
3. Magnetic Quantum Number ( ml )
Orientation of orbital
Specifies the exact orbitalwithin each sublevel
C. Quantum Numbers
px py pz
C. Quantum Numbers
Orbitals combine to form a spherical
shape.
2s
2pz2py
2px
C. Quantum Numbers
4. Spin Quantum Number ( ms )
Electron spin +½ or -½
An orbital can hold 2 electrons that spin in opposite directions.
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”:
Feeling overwhelmed?
Read Section 4-2!
IV. Electron Configuration(p. 105 - 116,
128 - 139)
Ch. 4 - Electrons in Atoms
A. General Rules
Pauli Exclusion Principle
Each orbital can hold TWO electrons
with opposite spins.
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 Electrons
Core Electrons
S 16e- [Ne] 3s2 3p4
1s2 2s2 2p6 3s2 3p4
B. Notation
Longhand Configuration
© 1998 by Harcourt Brace & Company
sp
d (n-1)
f (n-2)
1234567
67
C. Periodic Patterns
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
s-block
1st Period
1s11st column of s-block
C. Periodic Patterns
Example - Hydrogen
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.
[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
Ch. 5 - The Periodic Table
I. History
A. Mendeleev
Dmitri Mendeleev (1869, Russian) Organized elements
by increasing atomic mass.
Elements with similar properties were grouped together.
There were some discrepancies.
A. Mendeleev
Dmitri Mendeleev (1869, Russian) Predicted properties of undiscovered
elements.
B. Moseley
Henry Mosely (1913, British)
Organized elements by increasing atomic number.
Resolved discrepancies in Mendeleev’s arrangement.
II. Organization of theElements
Ch. 5 - The Periodic Table
MetalsNonmetalsMetalloids
A. Metallic Character
Main Group ElementsTransition MetalsInner Transition Metals
B. Blocks
III. Periodic Trends
Ch. 5 - The Periodic Table
0
50
100
150
200
250
0 5 10 15 20Atomic Number
Ato
mic
Ra
diu
s (
pm
)
A. Periodic Law
When elements are arranged in order of
increasing atomic #, elements with similar
properties appear at regular intervals.
0
50
100
150
200
250
0 5 10 15 20Atomic Number
Ato
mic
Ra
diu
s (
pm
)
B. Chemical Reactivity
Families Similar valence e- within a group result in
similar chemical properties
B. Chemical Reactivity
Alkali MetalsAlkaline Earth MetalsTransition MetalsHalogensNoble Gases
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
Atomic Radius
0
50
100
150
200
250
0 5 10 15 20Atomic Number
Ato
mic
Ra
diu
s (
pm
)
D. Atomic Radius
Li
ArNe
KNa
1
2
3
4 5
6
7
Atomic Radius Increases to the LEFT and DOWN
D. Atomic Radius
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
First Ionization Energy
0
500
1000
1500
2000
2500
0 5 10 15 20Atomic Number
1s
t Io
niz
ati
on
En
erg
y (k
J)
E. Ionization Energy
KNaLi
Ar
NeHe
1
2
3
4 5
6
7
First Ionization Energy Increases UP and to the RIGHT
E. Ionization Energy
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
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
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
1
2
3
4 5
6
7
Melting/Boiling Point Highest in the middle of a period.
F. Melting/Boiling Point
Ionic Radius
Cations (+)
lose e-
smaller
© 2002 Prentice-Hall, Inc.
Anions (–)
gain e-
larger
G. Ionic Radius
Which atom has the larger radius?
Be or Ba
Ca or Br
Ba
Ca
Examples
Which atom has the higher 1st I.E.?
N or Bi
Ba or Ne
N
Ne
Examples
Which atom has the higher melting/boiling point?
Li or C
Cr or Kr
C
Cr
Examples
Which particle has the larger radius?
S or S2-
Al or Al3+
S2-
Al
Examples
Periodic Trends Summary