Chapter 2 - 1
ISSUES TO ADDRESS...
• What promotes bonding?
• What types of bonds are there?
• What properties are inferred from bonding?
Chapter 2: Atomic Structure & Interatomic Bonding
Chapter 2 - 2
Atomic Structure (Freshman Chem.)
} 1.67 x 10-27 kg
Chapter 2 - 3
Atomic Structure (Freshman Chem.)
• atomic mass unit, amu ( a way to compute atomic weight) = 1/12 mass of 12C (most common isotope of carbon)
• In one mole of a substance, there are 6.022x1023 atoms or molecules. • 1 amu/atom(or molecule in a compound) = 1 g/mol Example: atomic weight of iron is 55.85 amu/atom, or 55.85 g/mol. • We use g/mol (or kg/kmol) in this book.
Chapter 2 - 4
Discuss Example Problem 2.1 in class. Four naturally occurring isotopes of Cerium: 0.185% of Ce136 with an atomic weight of 135.907 amu 0.251% of Ce138 with an atomic weight of 137.906 amu 88.45% of Ce140 with an atomic weight of 139.905 amu 11.114% of Ce142 with an atomic weight of 141.909 amu Calculate the average atomic weight of Ce.
The average atomic weight of a hypothetical element M, �̅�𝑀
�̅�𝑀 = �𝑓𝑖𝑀𝐴𝑖𝑀𝑖
𝑓𝑖𝑀: fraction of occurrence of isotope 𝑖 (i.e., the percentage-of- Occurrence divided by 100) 𝐴𝑖𝑀: the atomic weight of the isotope
Chapter 2 - 5
Atomic Structure • Some of the following properties
1) Electrical 2) Thermal 3) Optical 4) Deteriorative
are determined by electronic structure (behavior of electrons in atoms and crystalline solids) • Quantum mechanics: set of principles and laws that govern
systems of atomic and subatomic entities
Chapter 2 - 6
Bohr model vs wave-mechanical model • Bohr Model (Significant limitations):
Electrons are assumed to revolve around the atomic nucleus in discrete orbitals
• Wave-mechanical model: electron’s position is considered to be the probability of an electron’s being at various locations around the nucleus.
Chapter 2 - 7
Electronic Structure-Electron Energy States • Electrons have wavelike and particulate properties. • Two of the wavelike characteristics are
– 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) K, L, M, N, O (1, 2, 3, etc.) = second-azimuthal (orbitals) s, p, d, f (0, 1, 2, 3,…, n -1)
• have discrete energy states • tend to occupy lowest available energy state.
Electrons...
Chapter 2 - 8
Electron Energy States • Quantum # Designation
n = principal (energy level-shell) K, L, M, N, O (1, 2, 3, etc.) = second-azimuthal (orbitals) s, p, d, f (0, 1, 2, 3,…, n -1)
Chapter 2 - 9
Electronic Structure-Electron Energy States Notes: • The smaller the principal quantum number, the
lower the energy level; For example: energy of a 1s < energy of a 2s < energy of a 3s
• Within each shell, the energy of a subshell level increases with the value of the l quantum number.
For example: energy of a 3d > energy of a 3p> energy of a 3s.
• There may be overlap in energy of a state in one shell with states in an adjacent shell, which is especially true of d and f states.
For example: The energy of a 3d state is generally greater than that of a 4s.
Chapter 2 - 10
Electron Configurations • Valence electrons – those that occupy the
outermost shell • Filled shells are 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 - 11
• Stable: Valence (outermost) electron shell (normally just s and p) are completely filled.
• Most elements: Electron configuration not stable. SURVEY OF ELEMENTS
Electron configuration
(stable)
...
... 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)
Atomic #
18 ... 36
Element 1s 1 1 Hydrogen 1s 2 2 Helium 1s 2 2s 1 3 Lithium 1s 2 2s 2 4 Beryllium 1s 2 2s 2 2p 1 5 Boron 1s 2 2s 2 2p 2 6 Carbon
... 1s 2 2s 2 2p 6 (stable) 10 Neon 1s 2 2s 2 2p 6 3s 1 11 Sodium 1s 2 2s 2 2p 6 3s 2 12 Magnesium 1s 2 2s 2 2p 6 3s 2 3p 1 13 Aluminum
... Argon ... Krypton
Adapted from Table 2.2, Callister & Rethwisch 9e.
Chapter 2 - 12
Electronic Configurations ex: Fe - atomic # = 26
valence electrons
1s
2s 2p
K-shell n = 1
L-shell n = 2
3s 3p M-shell n = 3
3d
4s
4p 4d
Energy
N-shell n = 4
1s2 2s2 2p6 3s2 3p6 3d 6 4s2
Adapted from Fig. 2.6, Callister & Rethwisch 9e. (From K. M. Ralls, T. H. Courtney, and J. Wulff, Introduction to Materials Science and Engineering, p. 22. Copyright © 1976 by John Wiley & Sons, New York. Reprinted by permission of John Wiley & Sons, Inc.)
Chapter 2 - 13
The Periodic Table • Columns: Similar Valence Structure
Adapted from Fig. 2.8, Callister & Rethwisch 9e.
Electropositive elements: Readily give up electrons to become + ions.
Electronegative elements: Readily acquire electrons to become - ions.
give
up
1e-
give
up
2e-
give
up
3e-
iner
t gas
es
acce
pt 1
e-
acce
pt 2
e-
O
Se
Te
Po At
I
Br
He
Ne
Ar
Kr
Xe
Rn
F
Cl S
Li Be
H
Na Mg
Ba Cs
Ra Fr
Ca K Sc
Sr Rb Y
Chapter 2 - 14
Smaller electronegativity
Larger electronegativity
Electronegativity • Ranges from 0.9 to 4.1, • Large values: tendency to acquire electrons.
Chapter 2 - 15
• Forces between to adjacent atoms: FN = FA + FR (FN : net force, FA : attractive force, FR : repulsive force) • At the state of equilibrium: FA + FR = 0 • Sometimes it is more convenient to work with the potential
energies (E) between two atoms instead of forces (F).
EN = EA + ER • At the state of equilibrium EN is minimum. (Bonding Energy)
Bonding Forces and Energies
Chapter 2 - 16
• Bonding energy: represents the energy required to separate two atoms to an infinite separation • A number of material properties depend on bonding energy: melting temperature, phase at room temperature, coefficient of thermal expansion, stiffness • Three types of primary (chemical) bonds: ionic, covalent, and metallic • There are secondary (physical) forces in many solids. - They are weaker than primary ones.
Bonding Forces and Energies
Chapter 2 - 17
Ionic bond metal + nonmetal
donates accepts electrons electrons
Dissimilar electronegativities
ex: MgO Mg 1s2 2s2 2p6 3s2 O 1s2 2s2 2p4 [Ne] 3s2
Mg2+ 1s2 2s2 2p6 O2- 1s2 2s2 2p6 [Ne] [Ne]
Chapter 2 - 18
• Occurs between + and - ions. • Requires electron transfer. • Large difference in electronegativity required. • Example: NaCl
Ionic Bonding
Na (metal) unstable
Cl (nonmetal) unstable
electron
+ - Attraction
Na (cation) stable
Cl (anion) stable
Chapter 2 - 19
Ionic Bonding • Energy – minimum energy most stable
– Energy balance of attractive and repulsive terms
Attractive energy EA
Net energy EN
Repulsive energy ER
Interatomic separation r
r A
n r B EN = EA + ER = + -
Adapted from Fig. 2.10(b), Callister & Rethwisch 9e.
Eq. 2.10
≅ 8
Chapter 2 - 20
Ionic Bonding
Ionic bonding is typified by ceramic materials, which are characteristically hard and brittle and, furthermore, electrically and thermally insulative.
Chapter 2 - 21
• Predominant bonding in Ceramics Examples: Ionic Bonding
Give up electrons Acquire electrons
NaCl MgO CaF 2 CsCl
Chapter 2 - 23
Each H: has 1 valence e-, needs 1 more
Electronegativities are the same.
Fig. 2.12, Callister & Rethwisch 9e.
Covalent Bonding • similar electronegativity ∴ share electrons
• Example: H2
shared 1s electron from 2nd hydrogen atom
H
H2
shared 1s electron from 1st hydrogen atom
H
Chapter 2 - 24
Covalent Bonding
• Most covalently bonded materials are electrical insulators, or, in some cases, semiconductors.
• It is difficult to predict the mechanical properties of covalently bonded materials on the basis of their bonding characteristics. (some strong, others weak; some brittle, others ductile)
Chapter 2 - 25
Bond Hybridization • Carbon can form sp3 hybrid
orbitals
Fig. 2.13, Callister & Rethwisch 9e.
Fig. 2.14, Callister & Rethwisch 9e. (Adapted from J.E. Brady and F. Senese, Chemistry: Matter and Its Changes, 4th edition. Reprinted with permission of John Wiley and Sons, Inc.)
Chapter 2 - 26
Covalent Bonding: Carbon sp3 • Example: CH4
C: has 4 valence e-, needs 4 more
H: has 1 valence e-, needs 1 more
Electronegativities of C and H are comparable so electrons are shared in covalent bonds.
Fig. 2.15, Callister & Rethwisch 9e. (Adapted from J.E. Brady and F. Senese, Chemistry: Matter and Its Changes, 4th edition. Reprinted with permission of John Wiley and Sons, Inc.)
Chapter 2 - 27
Metallic Bonding • Metallic bonding is found in metals and their alloys. • Valence electrons are not bound to any particular atom in the
solid and are more or less free to drift throughout the entire metal (electron cloud).
• Metals are good conductors of both electricity and heat as a consequence of their free electrons.
Chapter 2 - 28
Arises from interaction between dipoles
• Permanent dipoles-molecule induced
• Fluctuating dipoles
-general case:
-ex: liquid HCl
-ex: polymer
Adapted from Fig. 2.20, Callister & Rethwisch 9e.
Adapted from Fig. 2.22, Callister & Rethwisch 9e.
Secondary Bonding (van der Waals)
asymmetric electron clouds
+ - + - secondary bonding
H H H H
H 2 H 2
secondary bonding
ex: liquid H 2
H Cl H Cl secondary bonding
secondary bonding + - + -
secondary bonding
Chapter 2 - 29
Mixed Bonding • covalent-ionic, covalent-metallic, metallic-ionic • Ionic-Covalent Mixed Bonding
% ionic character %IC =
where XA & XB are the electronegativities of elements
%) 100 ( x
Ex: MgO XMg = 1.3 XO = 3.5
Chapter 2 - 30
Type Ionic
Covalent
Metallic
Secondary
Bond Energy Large!
Variable large-Diamond small-Bismuth
Variable large-Tungsten small-Mercury
smallest
Comments ceramics
semiconductors, ceramics polymer chains
metals
inter-chain (polymer) inter-molecular
Summary: Bonding
Chapter 2 - 31
Ceramics (Ionic & covalent bonding):
Large bond energy large Tm large E small α
Metals (Metallic bonding):
Variable bond energy moderate Tm variable (moderate) E moderate α
Summary: Primary Bonds
Polymers (Covalent & Secondary):
Dominant Secondary bonding small Tm small E large α
Tm: melting temperature E: bonding energy α: thermal energy expansion
Chapter 1 - 32
2.1 2.2 2.4 2.9 2.10 2.17 2.18
Suggested Problems for Chapter 2
2.25 2.1FE 2.2FE 2.3FE 2.4FE