1 atomic spectra blackbody radiation is the visible glow that solid objects emit when heated. max...
TRANSCRIPT
1
Atomic SpectraAtomic Spectra
• Blackbody radiation is the visible glow that solid
objects emit when heated.
• Max Planck (1858–1947): proposed the energy is
only emitted in discrete packets called quanta.
• The amount of energy depends on the frequency:
E h
hc h 6.626 10 34 J s
2
Atomic SpectraAtomic Spectra
• Albert Einstein (1879–1955): • Used the idea of quanta to explain the photoelectric effect.
• He proposed that light behaves as a stream of particles called photons.
3
Atomic SpectraAtomic Spectra
• A photon’s energy must exceed a minimum threshold for electrons to be ejected.
• Energy of a photon depends only on the frequency.
4
Atomic SpectraAtomic Spectra
• For red light with a wavelength of about 630 nm,
what is the energy of a single photon and one mole
of photons?
E h
hc h 6.626 10 34 J s
5
Wave–Particle DualityWave–Particle Duality
• Louis de Broglie (1892–1987): Suggested waves
can behave as particles and particles can behave
as waves. This is called wave–particle duality.
For Light : h
mc
h
p
For a Particle : h
mv
h
p
6
Quantum MechanicsQuantum Mechanics
• Niels Bohr (1885–1962): Described atom as
electrons circling around a nucleus and concluded
that electrons have specific energy levels.
• Erwin Schrödinger (1887–1961): Proposed
quantum mechanical model of atom, which focuses
on wavelike properties of electrons.
7
Quantum MechanicsQuantum Mechanics
• Werner Heisenberg (1901–1976): Showed that it
is impossible to know (or measure) precisely both
the position and velocity (or the momentum) at the
same time.
• The simple act of “seeing” an electron would
change its energy and therefore its position.
8
Quantum MechanicsQuantum Mechanics
)()4()( :position selectron'in y Uncertaint
4))(( :Principlety UncertainHeisenberg
m
hx
hmx
9
Quantum MechanicsQuantum Mechanics
• Erwin Schrödinger (1887–1961): Developed a
compromise which calculates both the energy of an
electron and the probability of finding an electron at any
point in the molecule.
• This is accomplished by solving the Schrödinger
equation, resulting in the wave function, .
10
Quantum NumbersQuantum Numbers
• Wave functions describe the behavior of electrons.
• Each wave function contains three variables called
quantum numbers:
• Principal Quantum Number (n)
• Angular-Momentum Quantum Number (l)
• Magnetic Quantum Number (ml)
11
Quantum NumbersQuantum Numbers
• Principal Quantum Number (n): Defines the size
and energy level of the orbital. n = 1, 2, 3,
• As n increases, the electrons get farther from the
nucleus.
• As n increases, the electrons’ energy increases.
• Each value of n is generally called a shell.
12
Quantum NumbersQuantum Numbers
• Angular-Momentum Quantum Number (l): Defines the three-dimensional shape of the orbital.
• For an orbital of principal quantum number n, the value of l can have an integer value from 0 to n – 1.
• This gives the subshell notation:
l = 0 = s orbital l = 1 = p orbital
l = 2 = d orbital l = 3 = f orbital
l = 4 = g orbital
13
Quantum NumbersQuantum Numbers
• Magnetic Quantum Number (ml): Defines the spatial orientation of the orbital.
• For orbital of angular-momentum quantum number, l, the value of ml has integer values from –l to +l.
• This gives a spatial orientation of:
l = 0 giving ml = 0
l = 1 giving ml = –1, 0, +1
l = 2 giving ml = –2, –1, 0, 1, 2, and so on…...
14
Quantum NumbersQuantum Numbers
• Spin Quantum Number:
• The Pauli Exclusion
Principle states that no
two electrons can have
the same four quantum
numbers.
15
Quantum NumbersQuantum Numbers
16
Electron Radial DistributionElectron Radial Distribution
17
Electron Radial DistributionElectron Radial Distribution
• s Orbital Shapes:
18
Electron Radial DistributionElectron Radial Distribution
• p Orbital Shapes:
19
Electron Radial DistributionElectron Radial Distribution
• d and f Orbital Shapes:
20
Effective Nuclear ChargeEffective Nuclear Charge
• Electron shielding leads to energy differences among orbitals within a shell.
• Net nuclear charge felt by an electron is called the effective nuclear charge (Zeff).
21
Effective Nuclear ChargeEffective Nuclear Charge
• Zeff is lower than actual nuclear charge.
• Zeff increases toward nucleus ns > np > nd > nf
• This explains certain periodic changes observed.
22
Effective Nuclear ChargeEffective Nuclear Charge
23
Electron Configuration of AtomsElectron Configuration of Atoms
• Pauli Exclusion Principle: No two electrons in an
atom can have the same quantum numbers (n, l,
ml, ms).
• Hund’s Rule: When filling orbitals in the same
subshell, maximize the number of parallel spins.
24
Electron Configuration of AtomsElectron Configuration of Atoms
• Rules of Aufbau Principle:
1. Lower n orbitals fill first.
2. Each orbital holds
two electrons; each
with different ms.
3. Half-fill degenerate
orbitals before pairing
electrons.
25
Electron Configuration of AtomsElectron Configuration of Atoms
Assigning Electrons to Atomic Orbitals
1. The number of electrons in an atom is equal to the atomic
number.
2. Assign electrons to the lowest energy orbitals first, then
build up.
26
Electron Configuration of AtomsElectron Configuration of Atoms
Assigning Electrons to Atomic Orbitals
3. No more than 2 electrons can occupy a single orbital: their
spins must be paired.
4. If more than one orbital is available at the same energy, add
single electrons with the same spin to each orbital before adding
two electrons to one orbital.
5. Use the periodic table as a guide.
27
Electron Configuration of AtomsElectron Configuration of Atoms
Writing Electron Configurations
Name the occupied atomic orbitals in the atom with the
number of electrons in each orbital written as a superscript.
Li: 1s22s1 Na: 1s22s22p63s1
Fe: 1s22s22p63s23p64s23d6
28
Electron Configuration of AtomsElectron Configuration of Atoms
Writing Electron Configurations
One may also write the configuation as a noble gas closed
shell plus the valence electrons present in the atom.
Li: [He]2s1 Na: [Ne]3s1
Fe: [Ar]4s23d6
29
Electron Configuration of AtomsElectron Configuration of Atoms
1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p
Increasing Energy
[He][Ne] [Ar] [Kr] [Xe] [Rn]
Core
30
Electron Configuration of AtomsElectron Configuration of Atoms
Li 1s2 2s1
1s 2s
Be 1s2 2s2
1s 2s
B 1s2 2s2 2p1
1s 2s 2px 2py 2pz
C 1s2 2s2 2p2
1s 2s 2px 2py 2pz
31
Electron Configuration of AtomsElectron Configuration of Atoms
N 1s2 2s2 2p3
1s 2s 2px 2py 2pz
O 1s2 2s2 2p4
1s 2s 2px 2py 2pz
Ne 1s2 2s2 2p5
1s 2s 2px 2py 2pz
S [Ne] [Ne] 3s2 3p4
3s 3px 3py 3pz
32
Electron Configuration of AtomsElectron Configuration of Atoms
33
Electron Configuration of AtomsElectron Configuration of Atoms
34
Electron Configuration of AtomsElectron Configuration of Atoms
• Anomalous Electron Configurations: Result from unusual stability of half-filled & full-filled subshells.
• Chromium should be [Ar] 4s2 3d4, but is [Ar] 4s1 3d5
• Copper should be [Ar] 4s2 3d9, but is [Ar] 4s1 3d10
• In the second transition series this is even more
pronounced, with Nb, Mo, Ru, Rh, Pd, and Ag having
anomalous configurations (Figure 5.20).
35
Periodic PropertiesPeriodic Properties
36
Electron Configuration of AtomsElectron Configuration of Atoms
Metallic Radius: One half of the distance between neighboring atoms in a solid sample.
Predicting Relative Atomic Radii:
1. The atom with the largest n is largest.
2. If n is equal, then the atom with the largest nuclear charge is smallest.
37
Atomic RadiiAtomic Radii
38
Atomic RadiiAtomic Radii
39
Atomic RadiiAtomic Radii