c h a p t e rc h a p t e r c h a p t e rc h a p t e r 5 5 periodicity & atomic structure...

C h a p t e rC h a p t e r 55Periodicity & Atomic StructurePeriodicity & Atomic Structure

Chemistry, 4th EditionMcMurry/Fay

Chemistry, 4th EditionMcMurry/Fay

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The Periodic TableThe Periodic Table

• The periodic table is the most important organizing

principle in chemistry.

• Chemical and physical properties of elements in

the same group are similar.

• All chemical and physical properties vary in a

periodic manner, hence the name periodic table.

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Electromagnetic RadiationElectromagnetic Radiation

Electromagnetic Radiation:

Energy propagated by an electromagnetic field. Electromagnetic radiation has both particle and wave nature.

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Electromagnetic RadiationElectromagnetic Radiation

Spectroscopy:

Branch of physical science that deals with the interaction of electromagnetic radiation with matter

Spectrometry:

The quantitative measurement of the intensity of radiation at a particular wavelength of light.

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Wave-Like Nature of LightWave-Like Nature of Light

Frequency (, Greek nu): Number of peaks that pass a given point per unit time.

Wavelength (, Greek lambda): Distance from one wave peak to the next.

Amplitude: Height measured from the center of the wave. The square of the amplitude gives intensity.

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• Speed of a wave is the wavelength (in meters)

multiplied by its frequency in reciprocal seconds.

Wavelength x Frequency = Speed

(m) x (s–1) = c (m/s–1)

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Particle-Like Nature of LightParticle-Like Nature of Light

Electromagnetic radiation can be described as a

stream of tiny particles, called photons, with a very

small mass and a very large velocity.

The velocity of photons traveling in a vacuum is:

c = 3.00 x 108 m/s

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Particle-Like Nature of LightParticle-Like Nature of Light

Where does a photon come from?

One photon is emitted when one atom or molecule in

an excited state relaxes to the ground state via the

emission of radiation.

E = h ν

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Atomic SpectraAtomic Spectra

• Atomic spectra:

Result from excited

atoms emitting light.

• Line spectra: Result

from electron

transitions between

specific energy levels.

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

4))(( :Principlety UncertainHeisenberg

m

hx

hmx

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• 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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• 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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• 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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• 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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• Spin Quantum Number:

• The Pauli Exclusion

Principle states that no

two electrons can have

the same four quantum

numbers.

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Electron Radial DistributionElectron Radial Distribution

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• s Orbital Shapes:

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• p Orbital Shapes:

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• 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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• 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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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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• 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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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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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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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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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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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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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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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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• 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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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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c h a p t e rc h a p t e r c h a p t e rc h a p t e r 5 5 periodicity & atomic structure...

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