Chapter 2: Atomic Structure &
Interatomic Bonding
ISSUES TO ADDRESS...
• What promotes bonding?
• What types of bonds are there?
• What properties are inferred from bonding?
Chapter 2 - 1
Atomic Structure (Freshman Chem.)
• atom –
electrons – 9.11 x 10-31 kg
protons
1.67 x 10-27 kg
neutrons
}
• atomic number = # of protons in nucleus of atom
= # of electrons of neutral species
• A [=] atomic mass unit = amu = 1/12 mass of 12C
Atomic wt = wt of 6.022 x 1023 molecules or atoms
1 amu/atom = 1g/mol
C
H
12.011
1.008 etc.
Chapter 2 - 2
AVAGADRO’S NUMBER = 6.022 x 1023 = NA
ATOMIC OR MOLECULAR WEIGHT =
NA x WEIGHT PER ATOM.
number of neutrons = N
number of protons = Z
A= Z + N
(2.1)
Chapter 2 - 3
Atomic Structure
• Valence electrons determine all of the
following properties
1)
2)
3)
4)
Chemical
Electrical
Thermal
Optical
Chapter 2 - 4
BOHR ATOM
Chapter 2 - 5
WAVE MECHANICAL MODEL OF
ATOM
Chapter 2 - 6
Electronic Structure
• Electrons have wavelike and particulate
properties.
– This means that electrons are in orbitals defined by a
probability.
– Each orbital at discrete energy level is determined by
quantum numbers.
Quantum #
Designation
n = principal (energy level-shell)
l = subsidiary (orbitals)
ml = magnetic
K, L, M, N, O (1, 2, 3, etc.)
s, p, d, f (0, 1, 2, 3,…, n-1)
1, 3, 5, 7 (-l to +l)
ms = spin
½, -½
Chapter 2 - 7
Electron Energy States
Electrons...
• have discrete energy states
• tend to occupy lowest available energy state.
4d
4p
N-shell n = 4
3d
4s
Energy
3p
3s
M-shell n = 3
Adapted from Fig. 2.4,
Callister & Rethwisch 8e.
2p
2s
L-shell n = 2
1s
K-shell n = 1
Chapter 2 - 8
SURVEY OF ELEMENTS
• Most elements: Electron configuration not stable.
Element
Hydrogen
Helium
Lithium
Beryllium
Boron
Carbon
...
Atomic #
1
2
3
4
5
6
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
...
Adapted from Table 2.2,
Callister & Rethwisch 8e.
Neon
Sodium
Magnesium
Aluminum
...
10
11
12
13
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
...
Argon
...
Krypton
18
...
36
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)
• Why? Valence (outer) shell usually not filled completely.
Chapter 2 - 9
Electron Configurations
• Valence electrons – those in unfilled shells
• Filled shells more stable
• Valence electrons are most available for
bonding and tend to control the chemical
properties
– example: C (atomic number = 6)
1s2 2s2 2p2
valence electrons
Chapter 2 - 10
Electronic Configurations
ex: Fe - atomic # = 26 1s2 2s2 2p6 3s2 3p6 3d 6 4s2
4d
4p
N-shell n = 4 valence
electrons
3d
4s
Energy
3p
3s
M-shell n = 3
Adapted from Fig. 2.4,
Callister & Rethwisch 8e.
2p
2s
L-shell n = 2
1s
K-shell n = 1
Chapter 2 - 11
give up 1egive up 2egive up 3e-
• Columns: Similar Valence Structure
accept 2eaccept 1einert gases
The Periodic Table
H
He
Li Be
O
F Ne
Na Mg
S
Cl Ar
K Ca Sc
Rb Sr
Y
Cs Ba
Se Br Kr
Te
I
Adapted from
Fig. 2.6,
Callister &
Rethwisch 8e.
Xe
Po At Rn
Fr Ra
Electropositive elements:
Readily give up electrons
to become + ions.
Electronegative elements:
Readily acquire electrons
to become - ions.
Chapter 2 - 12
Electronegativity
• Ranges from 0.7 to 4.0,
• Large values: tendency to acquire electrons.
Smaller electronegativity
Larger electronegativity
Adapted from Fig. 2.7, Callister & Rethwisch 8e. (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.
Chapter 2 - 13
Ionic bond – metal
+
donates
electrons
nonmetal
accepts
electrons
Dissimilar electronegativities
ex: MgO
Mg
1s2 2s2 2p6 3s2
[Ne] 3s2
Mg2+ 1s2 2s2 2p6
[Ne]
O
1s2 2s2 2p4
O2- 1s2 2s2 2p6
[Ne]
Chapter 2 - 14
Electrons in different shells
Chapter 2 - 15
Electrons in Sodium and Chlorine
TABLE 2.2 / P 25
3s1
3s2 3p5
Chapter 2 - 16
•
•
•
•
Ionic Bonding
Occurs between + and - ions.
Requires electron transfer.
Large difference in electronegativity required.
Example: NaCl
Na (metal)
unstable
Cl (nonmetal)
unstable
electron
Na (cation)
stable
-
+
Coulombic
Attraction
Cl (anion)
stable
Chapter 2 - 17
Chapter 2 - 18
FORCES AND ENERGIES
Chapter 2 - 19
Chapter 2 - 20
Chapter 2 - 21
Bonding Forces and Energies
2.13 Calculate the force of attraction
between a K+ and an O2- ion the
centers of which are separated by a
distance of r0 =1.5 nm.
Solution
The attractive force between two ions
FA is just the derivative with respect to
the interatomic separation of the
attractive energy expression, Equation
2.8, which is just
Chapter 2 - 22
FA =
dE A
dr
=
 A 
d  
 r 
dr
=
A
r
2
The constant A in this expression is
defined in footnote 3. Since the valences
of the K+ and
O2- ions
(Z1 and Z2) are +1 and -2, respectively,
Z1 = 1 and Z2 = 2, then
Chapter 2 - 23
FA =
=
(Z 1 e) (Z 2 e)
4  0 r 2
(1 )( 2 ) (1 .602  10  19 C ) 2
(4 )(  ) (8.85  10  12 F/m ) (1 .5  10  9 m) 2

=2.05  10^(-10 ) N
Chapter 2 - 24
IONIC FORCE / P 31 FOOT-NOTE
F= (Z1 *Z2 * e^2)/(4*π*ε0*r^2);
e= 1.602 *10^(-19) COULOMBS ;
ε0 = 8.85 * 10^(-12 )
Z1, Z2 = VALENCIES OF IONS
Chapter 2 - 25
Ionic Bonding
• Energy – minimum energy most stable
– Energy balance of attractive and repulsive terms
EN = EA + ER =

A
r
+
B
rn
Repulsive energy ER
Interatomic separation r
Net energy EN
Adapted from Fig. 2.8(b),
Callister & Rethwisch 8e.
Attractive energy EA
Chapter 2 - 26
Examples: Ionic Bonding
• Predominant bonding in Ceramics
NaCl
MgO
CaF 2
CsCl
Give up electrons
Acquire electrons
Adapted from Fig. 2.7, Callister & Rethwisch 8e. (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.
Chapter 2 - 27
Covalent Bonding
• similar electronegativity  share electrons
• bonds determined by valence – s & p orbitals
dominate bonding
• Example: CH4
C: has 4 valence e-,
needs 4 more
H: has 1 valence e-,
needs 1 more
Electronegativities
are comparable.
H
CH 4
H
C
H
shared electrons
from carbon atom
H
shared electrons
from hydrogen
atoms
Adapted from Fig. 2.10, Callister & Rethwisch 8e.
Chapter 2 - 28
Primary Bonding
• Metallic Bond -- delocalized as electron cloud
• Ionic-Covalent Mixed Bonding
% ionic character =
2

(X A X B )


4
1  e




 x (100%)


where XA & XB are Pauling electronegativities
Ex: MgO
% ionic character
XMg = 1.2
XO = 3.5
2
( 3 .5 1 .2 )



4
 1 e




x (100%)  73.4% ionic


Chapter 2 - 29
METALLIC BONDING
Chapter 2 - 30
SECONDARY BONDING
Arises from interaction between dipoles
• Fluctuating dipoles
asymmetric electron
clouds
+
-
+
secondary
bonding
-
ex: liquid H 2
H2
H2
H H
H H
secondary
bonding
Adapted from Fig. 2.13,
Callister & Rethwisch 8e.
• Permanent dipoles-molecule induced
-general case:
-ex: liquid HCl
-ex: polymer
+
-
H Cl
secondary
bonding
+
secondary
bonding
H Cl
Adapted from Fig. 2.15,
Callister & Rethwisch 8e.
secondary bonding
Chapter 2 - 31
Summary: Bonding
Comments
Type
Bond Energy
Ionic
Large!
Nondirectional (ceramics)
Covalent
Variable
large-Diamond
small-Bismuth
Directional
(semiconductors, ceramics
polymer chains)
Metallic
Variable
large-Tungsten
small-Mercury
Nondirectional (metals)
Secondary
smallest
Directional
inter-chain (polymer)
inter-molecular
Chapter 2 - 32
Properties From Bonding: Tm
• Bond length, r
• Melting Temperature, Tm
Energy
r
• Bond energy, Eo
ro
Energy
r
smaller Tm
unstretched length
ro
r
Eo =
“bond energy”
larger Tm
Tm is larger if Eo is larger.
Chapter 2 - 33
Properties From Bonding : a
• Coefficient of thermal expansion, a
length, L o
coeff. thermal expansion
unheated, T1
DL
= a(T2 -T1)
Lo
DL
heated, T 2
• a ~ symmetric at ro
Energy
unstretched length
ro
E
o
E
o
r
a is larger if Eo is smaller.
larger a
smaller a
Chapter 2 - 34
Summary: Primary Bonds
Ceramics
(Ionic & covalent bonding):
Metals
(Metallic bonding):
Polymers
(Covalent & Secondary):
Large bond energy
large Tm
large E
small a
Variable bond energy
moderate Tm
moderate E
moderate a
Directional Properties
Secondary bonding dominates
small Tm
small E
large a
Chapter 2 - 35
ANNOUNCEMENTS
Reading:
Core Problems:
Self-help Problems:
Chapter 2 - 36