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CHEM 1311A
E. Kent Barefield
Course web page
http://web.chemistry.gatech.edu/~barefield/1311/chem1311a.html
Two requests: cell phones to silent/off
no lap tops in operation during class
Bring your transmitter to class on Friday
Outline for first exam period
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Atomic structure and periodic properties
Structures and bonding models for covalent compounds of p-block
elements
–
Lewis Structures, Valence Shell Electron Repulsion (VSEPR)
concepts and oxidation state
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Hybridization
Composition and bond energies in binary p-block compounds
Structure and bonding in organic compounds
First Exam – Wednesday, September 10
Atomic structure and periodic properties
•
Early experiments concerning atomic structure and
properties of electromagnetic radiation
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Bohr model of the hydrogen atom
Demise of Bohr model
Wave equation and wave model for atom
Wave functions and properties
Orbital representations
Multielectron atoms
Periodic properties and their origin
Properties of Atoms
•
Consist of small positively charged nucleus surrounded by
negatively charged electrons, some at large distances from
the nucleus
•
Nucleus consists of positively charged protons and neutral
neutrons
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Charges of proton and electron are equal
•
The mass number of an atom is equal to the number of
protons plus the number of neutrons
•
Isotopes are atoms with the same atomic number but
different mass numbers, i.e., 1H, 2H, 3H or 12C, 13C, 14C
The atomic number (and nuclear charge) of an atom is
equal to the number of protons in its nucleus.
Atomic Spectra
Irradiation from electronically excited atoms is not continuous
Visible region line spectra could be fitted to simple (empirical) formula
1/λ = 1.097 x 107 m-1 (1/22-1/n2)
E = hc/λ = 2.178 x 10-18 J (1/22-1/n2)
Later work showed several series of lines
E = 2.178 x 10-18 J (1/n22-1/n12 ) (n2<n1)
Spectrum of Electromagnetic Radiation
Region
Wavelength
(Angstroms)
Wavelength
(centimeters)
Frequency
(Hz)
Energy
(eV)
Radio
> 109
> 10
< 3 x 109
< 10-5
Microwave
109 - 106
10 - 0.01
3 x 109 - 3 x 1012
10-5 - 0.01
Infrared
106 - 7000
0.01 - 7 x 10-5
3 x 1012 - 4.3 x 1014
0.01 - 2
Visible
7000 - 4000
Ultraviolet
4000 - 10
4 x 10-5 - 10-7
7.5 x 1014 - 3 x 1017
3 - 103
X-Rays
10 - 0.1
10-7 - 10-9
3 x 1017 - 3 x 1019
103 - 105
γ-Rays
< 0.1
< 10-9
> 3 x 1019
> 105
7 x 10-5 - 4 x 10-5 4.3 x 1014 - 7.5 x 1014
2-3
Visible region of electromagnetic spectrum
400 nm
4000 Å
4x10-5 cm
Blue green
Green
Yellow green
Yellow
Orange
Red
Purple
Violet
Indigo
Blue
500
600
Wavelength, nm
Red
Blue
Orange
400
Blue green
Indigo
Yellow
Complementary color
Violet
Color of transmitted light
Yellow green
700 nm
7000 Å
7x10-5 cm
700
800
Properties of Electromagnetic Radiation
c = λν
c = 2.998 x 108 m s-1 (vacuum); λ in meters; ν in cycles per sec or s-I
1 cycle per sec is 1 Hz (Hertz)
E = hν = hc/λ
With Einstein's relationship E = mc2
hc/λ = mc2
λ= h/mc
Bohr Model
Rule 1. Electron can exist in stationary states; requires fixed energy.
Rule 2. Possible stationary states determined by quantization of angular
momentum (mevr) in units of h/2π (mevr = nh/2π)
Rule 3. Transitions between stationary states occur with emission or
absorption of a quantum of energy, ΔE = hν
For a stationary state to occur for the electron moving in a planetary orbital
about the nucleus, the centripetal force must equal the electrostatic attraction
of the electron by the nucleus
or
Employing the expression for angular momentum from Rule 2:
Note that radius of orbits increase with n2 and decrease with Z
Total Energy for atom (E) = Kinetic Energy (KE)+ Potential Energy (PE)
from the previous slide
so
Substituting for r :
E = -(2.18 x 10-18 J) Z2/n2
E = -(13.6 eV) Z2/n2
Constants
me = 9.10939 x 10-11 kg
e = 1.60218 x 10-19 C
c = 8.8538 x 10-12 C2 m-1 J-1
h = 6.62608 x 10-34 J s
1 eV = 1.6022 x 10-19 J
ΔE = final energy – initial energy
ΔE = -13.6 Z2 eV (1/nf2 - 1/ni2)
Further evidence for unique properties of light
Photoelectric Effect
K.E.(ejected electrons) = hν(impinging light) - w(work function)
Compton Effect
Momentum and energy conserved in collisions of light (photons) with electrons
Demise of Bohr Model - Introduction of Wave Model
Heisenberg Indeterminacy (Uncertainty) Principle
Δx•Δ(mv) ≥ h/2π
Exact position and exact momentum cannot both be known simultaneously
De Broglie's Hypothesis
λ = h/mv
Particles have wave lengths; proof came from diffraction of electron beam
Schroedinger's Wave Equation
∂2Ψ/∂x2 + ∂2Ψ/∂y2 + ∂2Ψ/∂z2 + 8π2m/h2(E-V)Ψ = 0
(HΨ=EΨ)
Solution of this equation requires conversion to spherical polar coordinates and the
introduction of three constants, which are what we know as quantum numbers, n, ℓ, mℓ
Interpretation of Schrödinger equation solutions
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Each combination of n, ℓ, and mℓ, value corresponds to an orbital
n values relate to the energy and size of the orbitals. n = 1, 2, 3···
ℓ values specify the total angular momentum of the electron and
determines the angular shape of the orbital. ℓ’s have letter
equivalents: ℓ = 0 (s), 1 (p), 2 (d), 3 (f), 4 (g), etc. ℓ = 0 ··· n-1
•
mℓ, specifies the orientation of the orbitals in space and the
angular momentum with respect to direction. mℓ = -ℓ··· 0 ··· +ℓ
•
With introduction of the fourth quantum number, ms = +1/2, -1/2
and the Pauli principle, which states that no two eIectrons in a
given atom can have the same four quantum numbers, the
development of electron configurations for multielectron atoms is
straight-forward.
•
E = -me4Z2/8εo2n2h2 = -13.6 eV Z2/n2 Same as for Bohr model!!!
Wave Functions
⎡ 4Z 3 (n − l − 1)! ⎤
Rn,l (r ) = ⎢ 4 3
3⎥
⎢⎣ n ao [( n + l )! ] ⎥⎦
Ψn,ℓ,m (r,θ,φ) = Rn,ℓ(r)·Θℓ,m ·Φm
ℓ
1/ 2
⎛ 2Zr
⎜⎜
⎝ nao
ℓ
ℓ
⎞ − Zr / nao 2l +1
⎟⎟e
Ln + l
⎠
R1,0(r) = 2Z3/2ao-3/2e-Zr/a
o
R2,0(r)= (1/2√2)Z3/2ao-3/2(2-Zr/ao)e-Zr/2a
o
R2,1(r) = (1/2√6)Z3/2ao-3/2(Zr/ao)e-Zr/2a
o
R3,0(r) = (1/4√3)Z3/2ao-3/2(27-18Zr/ao+2Z2r2/9ao2)e-Zr/3a
o
ℓ
mℓ
Θ
φ
Θℓ,mℓ(θ) = sinmℓ2·Pℓ(cos θ)
0
0
1/√2
1/√2π
Φm (φ) = (1/√2π)e±imℓφ
1
0
(√3/2)cos θ
1/√2π
1
±1
(√3/2)sin θ
1/√2πe±iφ
2
0
(√5/8)(3cos2θ-1)
1/√2π
ℓ
1s Radial functions
30
6
25
5
20
4
R1,0(r) = 2ao-3/2e-r/a
o
15
3
R1,02(r) = 4ao-3e-2r/a
o
10
2
r2R1,02(r) = 4ao-3r2e-2r/a
o
5
1
0
0
0
1
2
3
r, ao
4
5
Representation of the 1s orbital
2s Radial functions
3.5
3
2.5
2
1.5
1
0.5
0
-0.5
0.7
0.6
0.5
0.4
0.3
0.2
0.1
0
-0.1
R2,0(r)= (1/2√2)ao-3/2(2-r/ao)e-r/2a
o
R2,02(r)= (1/8)ao-3(4-2r/ao+r2/ao2)e-r/a
o
r2R2,02(r)= (1/8)ao-3(4-2r/ao+r2/ao2)r2e-r/a
o
0
5
10
r, ao
15
Representations of p orbitals
Θℓ,mℓ(θ) =
sinmℓ2·Pℓ(cos
Φm (φ) = (1/√2π)e±imℓφ
ℓ
θ)
ℓ
mℓ
Θ
φ
0
0
1/√2
1/√2π
1
0
(√3/2)cos θ
1/√2π
1
±1
(√3/2)sin θ
1/√2πe±iφ
2
0
(√5/8)(3cos2θ-1)
1/√2π
Representations of p orbitals
Representations of orbitals
Conical node in dz2 is a sum of two planes:
dz2 = dz2–x2 + dz2–y2
Representations of orbitals
(+)
(-)
(-)
(+)
(+)
(+)
(-)
(+)
(-)
(+)
(-)
(+)
(-)
(-)
(+)
(+)
(-)
(-)
(+)
Comparison of ns radial distribution
functions
1.2
1
1s
0.8
0.6
2s
0.4
3s
0.2
0
0
2
4
6
8
10
r, ao
12
14
16
18
Comparison of 1s, 2s and 2p radial
distribution functions
1.2
1
1s
0.8
0.6
2p
0.4
2s
0.2
0
0
2
4
6
r, ao
8
10
12
Angular functions
as contour maps
• Angular nodes
• Radial nodes
• Relative size
Contour plots of H 2p and C 2p
3
2
Contour lines are set at 0.316 of
maximum values of Ψ2(r,θ,φ)
1
0
-1.5
-1
-0.5
0
-1
-2
-3
0.5
1
1.5
Multi-electron atoms; Aufbau Principle
n=1
s
H
p
1
0.8
0.4
0.2
[He]
Be [He]
B
[He]
C
[He]
N
[He]
O [He]
F
s
1.2
0.6
He
Li
n=2
[He]
Ne [He]
0
0
2
4
6
r, ao
8
10
12
More on multi-electron atoms
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Electron configurations in the third period (n=3) mirror those for
the second period
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In the fourth period (n=4) the 3d orbital energies are lower than
the 4p so that these orbitals fill before the 4p. The elements
with d electrons are referred to as transition elements.
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Electron configurations of the first transition element series are
4s23dn except for Cr and Cu, which are 4s13d5 and 4s13d10,
respectively. These irregularities are a result of lower electronelectron repulsion energies.
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All transition metal cations with + 2 or greater charge have dn
electron configurations.
Mn: [Ar]4s23d5
Mn2+: [Ar]4s03d5
Ti:
Ti3+:
[Ar]4s23d2
Ni: [Ar]4s23d8
[Ar]4s03d1
Ni2+: [Ar]4s03d8
The Periodic Table and the Aufbau Principle
(n)s
(n)p
n =1
2
3
4
5
6
7
(n–2)f
(n–1)d
Periodic properties
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Ionization energy
Atomic and ionic radii
Electron Affinity
Electronegativity
Ionization energy
Ionization energy is the minimum energy required to remove an
electron from a gaseous atom (or other species). E = E+ + e2500
1500
1000
2500
500
2000
0
0
20
40
60
Z
80
IE/eV
IE/kJ mol-1
2000
1500
1000
500
0
0
4
8
12
Z
16
Effective nuclear charge
s & p orbital energies vs Z
0
-10
s
p
E/eV
-20
Zeff = Z – shielding constant
-30
-40
-50
-60
Li
Be
B
C
N
O
F
Effective nuclear charges
1s
2s
2p
H
1.0
Li
1.28
He
1.69
Be
1.91
B
2.58
2.42
C
3.22
3.14
N
3.85
3.83
O
4.49
4.45
F
5.13
5.10
Ne
5.76
5.76
Atomic and ionic radii
Atomic radii as function of Z
Radii of some common ions, Å
Li+
0.59
Be2+
0.27
N31.71
O21.40
F1.33
Na+
1.02
Mg2+
0.72
Al3+
0.53
P32.12
S21.84
Cl1.81
K+
1.38
Ca2+
1.00
Ga3+
0.62
As32.22
Se21.98
Br1.96
Rb+
1.49
Sr2+
1.16
In3+
0.79
Cs+
1.70
Ba2+
1.36
Tl3+
0.88
Fe2+ 0.75, Fe3+ 0.69
I2.20
Radial distribution functions for
alkali metal ions
Electron affinity
Electron affinity is the energy change associated with addition
of an electron to a atom or ion. E + e- = E400
EA/kJ mol-1
300
200
100
0
-100
-200
H
Li Be B
C
N
O
F
Na Mg Al
Si
P
S
Cl
K
Ca
Note that EA’s have been historically recorded as the negative of
the energy associated with the electron attachment reaction.
Electronegativity
Electronegativity is the ability of an atom to attract electron
density to itself in a chemical bond
Pauling electronegativities
H
2.20
Li
0.98
Be
1.57
B
2.04
C
2.55
N
3.04
O
3.44
F
3.98
Na
0.93
Mg
1.31
Al
1.61
Si
1.90
P
2.19
S
2.58
Cl
3.16
K
0.82
Ca
1.00
Ga
1.81
Ge
2.01
As
2.18
Se
2.55
Br
2.96
Rb
0.82
Sr
0.95
In
1.78
Sn
1.96
Sb
2.05
Te
2.10
I
2.66
Cs
0.79
Ba
0.89