chapter 2: atomic bonding. reading assignment ch. 2 and 3 of textbook

Chapter 2: Atomic bonding

Reading assignment

Ch. 2 and 3 of textbook

Homework No. 1

Problems 2-8, 2-9, 2-13.

Homework No. 2

Problems 3-13, 3-15, 3-17, 3-19, 3-20, 3-21, 3-27.

Effect of atomic bonding

Example: carbon exists as graphite (soft with greasy feeling) or diamond (hardest known material)

Atomic & electronic

configuration

Bonding

BOND STRENGTH

Mechanical & Physical

Properties

graphite diamond

Primary and secondary bonding primary bonds: strong atom-to-atom attractions produced by changes in electron position of the valence e– . Example : covalent atom between two hydrogen atoms

secondary bonds: much weaker. It is the attraction due to overall “electric fields”, often resulting from electron transfer in primary bonds. Example: intramolecular bond between H2 molecules gas

Electronic configurations

Valence electrons They represent the ability of an element to

enter into chemical combination with others. Valence es

− participate in the bonding between atoms.

Valence = # of electrons in outermost combined sp level.

Examples of the valence are:Mg: 1s22s22p63s2 valence = 2Al: 1s22s22p63s23p1 valence = 3 Ge: 1s22s22p63s23p63d104s24p2 valence = 4

Primary bonding types

Ionic bondingCovalent bondingMetallic bonding

Ionic bonding

Ionic bonding in NaCl

3s1

3p6

SodiumAtom

Na

ChlorineAtom

Cl

Sodium IonNa+

Chlorine IonCl -

IONIC

BOND

2-15

Ionic bondingIonic Bond:

The attractive bonding forces are coulombic (different polarities):

Interionic force

Force of attraction between Na+ and Cl- ions

Z1 = +1 for Na+, Z2 = -1 for Cl-

e = 1.60 x 10-19 C , ε0 = 8.85 x 10-12 C2/Nm2

a0 = Sum of Radii of Na+ and Cl- ions

= 0.095 nm + 0.181 nm = 2.76 x 10-10 m

NC

aeZZF attraction

910-212-

219

2

0

2

21 1002.3m) 10 x /Nm2)(2.76C 10 x 8.85(4

)1060.1)(1)(1(

4

+

2-18

Interatomic spacing The equilibrium distance between atoms is

caused by a balance between repulsive and attractive forces.

Equilibrium separation occurs when no net force acts to either attract or separate the atoms or the total energy of the pair of atoms is at a minimum.

For a solid metal the interatomic spacing is equal to the atomic diameter or 2r.

For ionically bonded materials, the spacing is the sum of the two different ionic radii.

Bond energyBonding force

and Energy curves for a Na+

& Cl- pair.

Since F = dE/da, the equilibrium bond length (ao) occurs @ F=0 and E is a minimum.

Coordination Coordination number (C.N.) = No. of nearest neighbors (radius R)

around (touching) a particular atom/ion (radius r).

C.N. depends on r/R ratio.

C.N. = 4 (in 2 dimensions)

Stable Stable Unstable(critical case)

r/R > min r/R = min r/R < min

C.N. = 2

Schematic drawing with nearest neighbors not in contact

C.N. = 3

Schematic drawing with nearest neighbors not in contact

C.N.= 43 dimensions

C.N. = 4 (3 dimensions)

C.N. = 8 (3 dimensions)

C.N. = 12 (3 dimensions)

C.N. = 63 dimensions

C.N. = 6 (3 dimensions)

C.N. = 4 (in 2 dimensions)C.N. = 6 (in 3 dimensions)

Stable Stable Unstable(critical case)

r/R > min r/R = min r/R < min

C.N. = 3

C.N. = 8 (3 dimensions)

C.N. = 4

Criteria of packing ions in a solid

Positive charge = negative charge Nearest neighbors of a cation are anions; nearest

neighbors of an anion are cations. (Nearest neighbors touch one another.)

The coordination number (CN) is determined by r/R, where r = radius of smaller ion (usually the cation), and R=radius of larger ion (usually the anion). The greater is r/R, the higher is CN.

The largest allowable CN is most favorable.

Summary on ionic bonding

The attractive force (energy) for two isolated ions is a function of distance.

Bonding is nondirectional.This is the predominant bonding type

in ceramics. 600-1500 kJ/mol (3-8 eV/atom)

bonding energies are large high Tm.

Covalent bonding

Directional bond due to the sharing of electrons

between atoms

Cl2 molecule

Planetary model

Actual electron density

Electron dot schematic

Bond line schematic

Example 1. Br2 (a bromine molecule)

Br has an outermost electronic configuration of 4s24p5, i.e.,

A single bond

A -bondEnd-to-end overlap

Example 2. O2 (an oxygen molecule)

O has an outermost electron configuration of 2s22p4, i.e.,

A double bond

A σ-bond (end-to-end overlap) together

with a π-bond (side-to-side overlap).

Double bond

Single bonds

Every carbon atom along the chain is four-fold coordinated.

The electronic configuration of carbon is 1s22s22p2, i.e.

An excited state of carbon with electronic configuration

sp3 hybridization

Mixing of an s electron cloud with three p

electron clouds

Methane molecule

Methane molecule

Four sp3 orbitals are directed symmetrically toward corners of regular tetrahedron.

This structure gives high hardness, high bonding strength (711KJ/mol) and high melting temperature (3550oC).

Carbon atom Tetrahedral arrangement in diamond

2-25

SiO4

tetrahedron in silicate glass

Silicon dioxide (SiO2)

Covalent network solids

Network of covalent bondsExamples: diamond, silicon,

etc.

Properties of covalent network solids

Materials have poor ductility. Poor electrical conductivity. Many ceramic, semiconductor and polymer materials

are fully or partly covalent.

Mixed bonding (ionic + covalent)

Few compounds exhibit pure ionic or pure

covalent bonding. the bond type degree

depends on their position in the Periodic Table. The greater the difference in

electronegativity, the more ionic is the bond.

Conversely, the smaller the difference, the larger is the degree of covalency.

Example of mixed ionic-covalent bonding: HF (a hydrogen fluoride molecule)

The electronic configuration of H is 1s1, i.e.,

There is one unpaired electron.

The electronic configuration of F is 1s22s22p5, i.e.

There is also one unpaired electron.

Metallic Bonding

Found in metallic elements (low electronegativities).

Give up their valence electrons to form a “sea or cloud” of electrons.

The valence electrons move freely within the electron sea and become associated with the ion cores. The free electrons shield the (+) charged ion cores from repulsion.

Atoms in metals are closely packed in crystal structure. Loosely bounded valence electrons are attracted towards nucleus

of other atoms. Electrons spread out among atoms forming electron clouds. These free electrons are reason for electric conductivity and ductility Since outer electrons are shared by many atoms, metallic bonds are Non-directional

Positive Ion

Valence electron charge cloud2-28

Metallic Bonding

Good thermal &

electrical

conductors

The free electrons in the “cloud” move freely under an applied voltage.

Bond energy

It is the energy required to create or break the bond.

Materials with high bond energy high strength and high melting point.

Ionic materials have a large bond energy due to the large difference in electronegativities between the ions.

Metals have lower bond energies, because the electronegativities of the atoms are similar.

Bond energy and melting temperature

Secondary bonding

Van der Waals bonding

Secondary bonding exists between virtually all Secondary bonding exists between virtually all molecules, but its presence is diminished if any molecules, but its presence is diminished if any primary bond is present. primary bond is present.

Dipoles are created when positive andnegative charge centers exist.

-q

Dipole moment=μ =q.d

q= Electric charged = separation distance

2-30

+q d

Skewed electron cloud

Secondary bonds are due to attractions of electric dipoles in atoms or molecules.

Electric dipole types

Permanent dipolesInduced dipoles

Permanent dipoles

Dipoles that do not fluctuate

with time

Hydrogen fluoride HF

Example of a permanent dipole

Permanent dipoles in water

Attraction between positive oxygen pole and negative hydrogen pole.

Water

105 0O

H

HDipole-dipole

interaction

2-33

Dipole-dipole interaction in water

Methane

Vector sum of four C-H dipoles is zero.

Symmetricalarrangement

of 4 C-H bondsCH4

No dipolemoment

CH3ClAsymmetrical

tetrahedralarrangement

Createsdipole

2-32

Methane

Methyl chloride

Cl- ions in green

Hydrogen bonding

Hydrogen bonds are dipole-dipole interaction between polar bonds

containing hydrogen atoms. It is a particularly strong type of secondary

bonding, due to the almost bare proton. Examples: water, HF, etc.

Induced dipolesNo permanent dipole momentStatistical fluctuation in electron

density distributionLondon dispersion forces (weak)Examples: argon (an inert gas),

methane (CH4 - a symmetric molecule)

Weak secondary bonds in noble gasses. Dipoles are created due to asymmetrical distribution of electron

charges. Electron cloud charge changes with time.

Symmetricaldistribution

of electron charge

Asymmetricaldistribution

(Changes with time)

2-31

London forces between methane molecules

Another example of mixed bonding types in a material

Graphite

Crystal forms of carbon

GraphiteDiamond Fullerene

Fullerene (a molecule)

Diamond

Graphite

The electronic configuration of carbon is 1s22s22p2, i.e.

An excited state of carbon with electronic configuration

sp2 hybridization

Mixing of an s electron cloud with

two p electron clouds

Bonding in graphite

In-plane bonding: covalent + metallic

Out-of-plane bonding: van der Waal’s bonding

Consequent properties of graphite

Van der Waal’s bonding between layers - ease of sliding between layers (application as lubricant)

Metallic bonding within a layer – high in-plane thermal and electrical conductivity

Bonding in benzene molecule

sp2 hybridization of the carbon atoms

Chemical composition of benzene is C6H6.

The carbon atoms are arranged in hexagonal ring. Single and double bonds alternate between the atoms.

CC

CC

C

CH

H

H

H

H

HStructure of benzene Simplified notations

2-27

Effect of atomic bonding on material properties

Modulus of elasticity

Related to the material stiffness.Defined as the amount that a material will

stretch when a force is applied.It is related to the slope of the force-

distance curve.

A steep dF/da slope gives a high modulus.

Coefficient of thermal expansion

Describes how much a material expands or contracts when its temperature changes.

Asymmetric energy trough resulting in thermal expansion phenomenon