1 Atomic Spectra Blackbody radiation is the visible glow that solid objects emit when heated. Max Planck (1858–1947): proposed the energy is only emitted.

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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

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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.

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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.

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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

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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

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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.

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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.

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Quantum MechanicsQuantum Mechanics

)()4()( :position selectron'in y Uncertaint

4))(( :Principlety UncertainHeisenberg

m

hx

hmx

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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, .

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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)

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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.

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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

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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…...

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Quantum NumbersQuantum Numbers

• Spin Quantum Number:

• The Pauli Exclusion

Principle states that no

two electrons can have

the same four quantum

numbers.

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Quantum NumbersQuantum Numbers

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Electron Radial DistributionElectron Radial Distribution

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Electron Radial DistributionElectron Radial Distribution

• s Orbital Shapes:

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Electron Radial DistributionElectron Radial Distribution

• p Orbital Shapes:

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Electron Radial DistributionElectron Radial Distribution

• d and f Orbital Shapes:

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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).

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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.

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Effective Nuclear ChargeEffective Nuclear Charge

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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.

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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.

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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.

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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.

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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

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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

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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

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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

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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

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Electron Configuration of AtomsElectron Configuration of Atoms

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Electron Configuration of AtomsElectron Configuration of Atoms

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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).

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Periodic PropertiesPeriodic Properties

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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.

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Atomic RadiiAtomic Radii

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Atomic RadiiAtomic Radii

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Atomic RadiiAtomic Radii