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