Transcript Chapter 5

Chapter 5
Periodicity and Atomic Structure
Development of the Periodic Table
• The periodic table is the most important
organizing principle in chemistry.
– Periodic table powerpoint – elements of a group have
similar properties
– Chapter 2 – elements in a group form similar formulas
– Predict the properties of an element by knowing the
properties of other elements in the group
Light and the Electromagnetic
Spectrum
• Radiation (light) composed waves of
energy
• Waves were continuous and spanned the
electromagnetic spectrum
Light and the Electromagnetic
Spectrum
Light and the Electromagnetic
Spectrum
Light and the Electromagnetic
Spectrum
• 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)
» C – speed of light - 2.9979 x 108 m/s–1
Electromagnetic Radiation and
Atomic Spectra
• Classical Physics does not explain
– Black-body radiation
– Photoelectric effect
– Atomic Line Spectra
Particlelike Properties of Electromagnetic
Radiation: The Plank Equation
• Blackbody radiation is the visible glow that solid
objects emit when heated.
• Max Planck (1858–1947): Developed a formula to
fit the observations. He proposed that 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
Particlelike Properties of Electromagnetic
Radiation: The Plank Equation
• A photon’s energy
must exceed a
minimum threshold
for electrons to be
ejected.
• Energy of a photon
depends only on
the frequency.
Electromagnetic Radiation and
Atomic Spectra
• Atomic spectra: Result from excited
atoms emitting light.
– Line spectra: Result from electron
transitions between specific energy
levels.
Electromagnetic Radiation and
Atomic Spectra
1/λ = R [1/m2 – 1/n2]
Quantum Mechanics and the
Heisenburg Uncertainty Principle
• 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.
Quantum Mechanics and the
Heisenburg Uncertainty Principle
• 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.
Wave Functions and Quantum
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, .
Wave Functions and Quantum
Mechanics
• 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)
Wave Functions and Quantum
Mechanics
• 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.
Wave Functions and Quantum
Mechanics
• 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
Wave Functions and Quantum
Mechanics
• 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…...
Wave Functions and Quantum
Mechanics
Problem
• Why can’t an electron have the following
quantum numbers?
– (a) n = 2, l = 2, ml = 1
(b) n = 3, l = 0, ml = 3
– (c) n = 5, l = –2, ml = 1
• Give orbital notations for electrons with the
following quantum numbers:
– (a) n = 2, l = 1, ml = 1
– (c) n = 3, l = 2, ml = –1
(b) n = 4, l = 3, ml = –2
The Shapes of Orbitals
• s Orbital Shapes:
The Shapes of Orbitals
• p Orbital Shapes:
The Shape of Orbitals
• d and f Orbital Shapes:
Orbital Energy Levels in
Multielectron Atoms
Orbital Energy Levels in
Multielectron Atoms
• Zeff is lower than actual
nuclear charge.
• Zeff increases
toward nucleus
ns > np > nd > nf
• This explains certain
periodic changes observed.
Orbital Energy Levels in
Multielectron Atoms
• 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).
Wave Functions and Quantum
Mechanics
• Spin Quantum
Number:
• The Pauli Exclusion
Principle states that no
two electrons can have
the same four quantum
numbers.x
Electron Configurations of
Multielectron 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.
Electron Configurations of
Multielectron 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.
Electron Configurations and
Multielectron Atoms

1s

2s
1s2 2s1
Be 
1s

2s
1s2 2s2
Li
B
  
1s 2s 2px 2py 2pz
1s2 2s2 2p1
C
  

1s 2s 2px 2py 2pz
1s2 2s2 2p2
Electron Configurations and
Multielectron Atoms
N
    
1s 2s 2px 2py 2pz
1s2 2s2 2p3
O
    
1s 2s 2px 2py 2pz
1s2 2s2 2p4
Ne     
1s 2s 2px 2py 2pz
1s2 2s2 2p5
[Ne]   
[Ne] 3s2 3p4
S

3s 3px 3py 3pz
Problems
• Give the ground-state electron configurations for:
– Ne (Z = 10)
Mn (Z = 25)
Zn (Z = 30)
– Eu (Z = 63)
W (Z = 74)
Lw (Z = 103)
• Identify elements with ground-state configurations:
– 1s2 2s2 2p4
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2 4d6
– 1s2 2s2 2p6
[Ar] 4s2 3d1
[Xe] 6s2 4f14 5d10 6p5
Electron Configurations and the
Periodic Table
Some Anomalous Electron
Configurations
• 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).
Electron Configurations and
Periodic Properties: Atomic Radii
Optional Homework
• Text – 5.24, 5.26, 5.28, 5.30, 5.32, 5.34,
5.44, 5.56, 5.58, 5.66, 5.68, 5.70, 5.72,
5.76, 5.78, 5.82, 5.84, 5.94, 5.98, 5.108
• Chapter 5 Homework online
Required Homework
• Assignment 5