Transcript Chapter 2: Atomic Structure & Interatomic Bonding

L02: Chapter 2: Atomic Structure
and Interatomic Bonding
• Much of the material in this chapter you already had in
freshman chemistry, although the terminology is different for
some things. It is assumed that you remember what you
learned in freshman chemistry.
• ISSUES ADDRESSED:
- What promotes bonding between atoms?
- The types of bonds in solids.
- All properties are influenced by bonding.
Last revised January 8, 2013 by W.R. Wilcox, Clarkson University.
Wavelike behavior of electrons
• We often use dots to represent individual electrons and lines for
pairs of electrons in a covalent bond.
• However, the electrons have a wavelike behavior such that we
can only give the probability of an electron being in a particular
volume.
• So a more accurate representation of an electron is as a cloud.
• Quantum mechanics is used to describe the wave behavior of
electrons, but is not covered in this course.
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Figure 2.3
(a) Bohr atom
(b) Wave mechanical
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Electronic structure of isolated atoms
• The characteristics below stem from their wavelike nature.
– electrons are in orbitals
– each orbital is at a discrete energy level determined by its quantum
numbers
– the letter designations below were given to bands observed in optical
emission and absorption, but not understood at the time.
Quantum Number
n = principal (energy level-shell)
l = angular (sub shell, shape)
ml = magnetic
Designation
1,2,3,4,5,6,7 (K, L, M, N, O,…)
s, p, d, f (n of them to max of 4)
- l to + l by integers, including 0
ms = spin
½, -½
 Dynamic periodic table
 Atomic orbitals
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Electron energy states of isolated atoms
Electrons have discrete energy states. They occupy the lowest
possible energy levels, unless excited by an external source of energy,
e.g. thermal energy or absorption of photons (in light).
4d
4p
N-shell n = 4
3d
4s
Energy
3p
3s
M-shell n = 3
2p
2s
L-shell n = 2
1s
K-shell n = 1
Two electrons of
opposite spin can
be in each level.
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Ground-state energy levels of some elements
Element
Atomic #
Hydrogen
1
Helium
2
Lithium
3
Beryllium
4
Boron
5
Carbon
6
...
Neon
10
Sodium
11
Magnesium
12
Aluminum
13
...
Electron configuration
1s 1
1s 2
(stable)
1s 2 2s 1
1s 2 2s 2
1s 2 2s 2 2p 1
1s 2 2s 2 2p 2
...
Argon
...
Krypton
1s 2 2s 2 2p 6 3s 2 3p 6
(stable)
...
1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 (stable)
18
...
36
1s 2 2s 2 2p 6
(stable)
1s 2 2s 2 2p 6 3s 1
1s 2 2s 2 2p 6 3s 2
1s 2 2s 2 2p 6 3s 2 3p 1
...
The electron configuration is stable only for the noble gases. Except
for noble gases, the outer shell is not completely filled and so one or
more electrons may be lost or gained to form an ion,
or shared in a covalent bond.
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Electronegativity
Table 2.7 in the text has the 1960 values of Pauling.
The revised values below are generally used now.
Those with high values tend to give up
one or more electrons.
Those with low values to accept
electrons.
Bonding in solids
• As two atoms approach one
another, they at first
experience an attraction.
• They repel one another when
they are brought very close.
• r0 is the equilibrium distance.
• The type of bonding in a solid
depends on the behavior of the
atoms' outer “valence”
electrons.
• Metallic: outer electrons shared in a cloud or sea.
• Ionic
– Cations have given up one or more electrons
– Anions have gained one or more electrons
• Covalent: atoms share outer electrons
• Mixed ionic and covalent
• Van der Waals: electrostatic due to non-uniform charge distribution. Weak
Typical ionic bond: metal + nonmetal
accepts
electrons
donates
electrons
• The greater the difference in electronegativity, the greater the tendency to
form an ionic bond.
• Consider magnesium and oxygen with electronegativities of 1.31 and 3.44.
• Here’s what happens when Mg and O come near one another:
1s2 2s2 2p6 3s2
(Ne + 3s2)
Mg2+ 1s2 2s2 2p6
(Ne)
cation
Mg
1s2 2s2 2p4
(Ne – 2p2)
O2- 1s2 2s2 2p6
(Ne)
anion
O
electron(s)
-
+
Coulombic
Attraction
Ionic bonding between one cation (+)
and one anion (-)
r0
Repulsive energy ER
Interatomic separation r
Net energy EN
EN = EA + ER =
-
A
r
+
Attractive energy EA
Stable at minimum energy E0 for radius r0.
Force = dE/dr = 0 at r0
r0
10
B
rn
Examples: Ionic Bonding
• Predominant bonding in Ceramics
NaCl
MgO
CaF 2
CsCl
Give up electrons
Acquire electrons
Adapted from Fig. 2.7, Callister & Rethwisch 4e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the
Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.
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Ionic bonding in a crystal
• In a crystal, a cation (+ charge) is attracted not only by the nearest
anions, but to a lesser extent by those farther away.
• Similarly, it is repelled by all other cations.
• The sum of the energy due to all attractions and repulsions is known
as the Madelung energy. This is approximately 60% greater than the
energy of attraction for isolated ions the same distance apart as in
the lattice.
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Covalent Chemical Bonds
• Atoms with almost the same electronegativity share electrons
leading to hybrid electronic structures.
• The bonds are very directional, unlike ionic bonds.
• Example:
H
CH 4
H
C
H
shared electrons
from carbon atom
H
shared electrons
from hydrogen
atoms
 Methane orbitals
• Hybrid orbitals
 Covalent orbitals
Mixed Ionic-Covalent Bonding
2
• Ionic-Covalent Mixed Bonding
(
X
X
)
A
B


4

• Approximate fraction ionic character  1 - e


where XA & XB are the two Pauling electronegativities.
Example: MgO. Using the 1960 values in the text,
XMg = 1.2 and XO = 3.5,
the equation above predicts that the bond between Mg and O has
about 73% ionic character and 27% covalent.
Using the revised values given on Wikipedia,
XMg = 1.31 and XO = 3.44,
the equation above predicts that the bond between Mg and O has
about 68% ionic character and 32% covalent.
For homework problems use the values in the text.


Metallic bonding
• Occurs with atoms that easily give up electrons.
• In a solid, these “conduction” electrons form a cloud or sea.
• No two electrons can have exactly the same quantum number, and
so they have a range of energies. Each “exists” throughout the solid.
• The attraction between the positively charged metal ions and the
electron cloud is what causes metallic bonding.
• Non directional.
• These “conduction” electrons
carry electric current and heat.
• Mixtures of metals sometimes
form intermetallic compounds.
• Animation (in full-screen
projection mode):
Another animation: http://mypchem.com/myp9/myp9c/myp9c_swf/metal_vib.htm
15
Secondary (van der Waals) bonds
Arises from interaction between electric dipoles
• Dipoles fluctuating rapidly and interacting
asymmetric electron
clouds
+
-
+
H2
-
H H
e.g. liquid H2
H2
H H
• Permanent dipoles
-general case:
+
-ex: liquid HCl
H Cl
-ex: polymer
-
+
-
H Cl
Within an organic molecule the bonding is
mostly covalent, while between molecules
the bonding is mostly van der Waals.
Hydrogen bonds
• Between hydrogen atoms and the nearby negative end of a
molecular dipole, to strongly electronegative atoms such as O or N.
• Partly covalent and partly electrostatic.
• Much stronger than van der Waals bonds.
• Determines the unusual properties of water liquid and solid.
• Also occurs with other molecules, and even between parts of
complex molecules such as proteins.
• http://en.wikipedia.org/wiki/Hydrogen_bond
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Table 2.3.
Properties From Bonding: Melting point
Interaction energy E
versus atomic separation r
Melting Temperature, Tm
Energy
Energy
unstretched bond length
ro
r
Eo
“bond energy”
ro
r
smaller Tm
larger Tm
Atomic separation r
Tm is larger when Eo is larger
r
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Properties From Bonding: Thermal expansion
Coefficient of thermal expansion, a
length, L o
coeff. thermal expansion
unheated, T1
DL
= a (T2 -T1)
Lo
DL
heated, T 2
Energy
ro
Eo
Eo
r
a is larger when Eo is smaller
larger a
smaller a
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