Crystal Binding (Bonding)Overview & Survey of Bonding Types Continued

Covalent Bondıng• Covalent Bonding takes place between atoms with small

differences in electronegativity which are close to each other in periodic table (between non-metals and non-metals).

• Covalent Bonds are formed by sharing of outer shell electrons (i.e., s & p electrons) between atoms rather than by electron transfer.

• This bonding can happen if the two atoms each share one of the other’s electrons, so the closed shell, noble gas valence electron configuration can be attained.

Covalent Bonding of 2 H Atoms The H2 Molecule

Interaction Potential

• Each electron in a shared pair is attracted to both nuclei involved in the bond. The approach, electron overlap, and attraction can be visualized as shown in the following (crude!) figure representing the nuclei & electrons in a hydrogen molecule.

e

e

Covalent network substances are brittle. If sufficient force is applied to a crystal, covalent bond are broken as the lattice is distorted. Shattering occurs rather than deformation of a shape.

Covalent network substances are brittle. If sufficient force is applied to a crystal, covalent bond are broken as the lattice is distorted. Shattering occurs rather than deformation of a shape.

BrittlenessBrittleness

They are hard because the atoms are strongly bound in the lattice, and are not easily displaced. They are hard because the atoms are strongly bound in the lattice, and are not easily displaced.

HardnessHardness

Poor conductors because electrons are held either on the atoms or within covalent bonds. They cannot move through the lattice.

Poor conductors because electrons are held either on the atoms or within covalent bonds. They cannot move through the lattice.

Electricalconductivity

Electricalconductivity

Very high melting points because each atom is bound by strong covalent bonds. Many covalent bonds must be broken if the solid is to be melted and a large amount of thermal energy is required for this.

Very high melting points because each atom is bound by strong covalent bonds. Many covalent bonds must be broken if the solid is to be melted and a large amount of thermal energy is required for this.

Melting point & boiling pointMelting point & boiling point

ExplanationExplanationPropertyProperty

Covalent Materials

Pictorial Comparison ofIonic & Covalent Bonding

In a Covalent Bond, neighboring atomsSHARE electrons in the bond.

• This sharing occurs with one or more pairs of electrons between the two atoms.

• Sometimes this sharing is not equal between the atoms. Molecules are formed.

• In solids, a covalent network (lattice) of atoms is formed. The number of covalent bonds formed is equal to the number of electrons necessary to achieve the

Noble gas valence electronicconfiguration for each atom.

Covalent Bonds

In an Ionic Bond, neighboring atomsTRANSFER electrons from one site to another.

• Positive & negative ions are formed.• In solids, an ionic solid or a salt is formed.

Example: NaCl Nao Na+ (cation) + e-

Clo + e Cl - (anion) Nao + Clo NaCl (salt)

Ionic Bonds(Contrast to Covalent Bonds)

Mixed Covalent & Ionic BondingNOTE!!

• Some bonding in many materialsCannot be classified as

pure covalent or pure ionic.• Instead,

Many bonds are best considered as mixtures of covalent & ionic bonds

or as having some characteristics of both covalent bonding & ionic bonding.

Bond Polarity• The Polarity of a bond is a measure of how unequal

(or not) the atoms are at electron sharing. It is another manifestation of The Relationship Between Electronegativity Difference and Bond Type.

Schematically, this can be viewed as follows: Bond Type Electronegativity Difference in Bonding Atoms

Nonpolar Covalent Bond Equal Sharing

Polar Covalent Bond Unequal Sharing

Ionic Bond Complete Electron Transfer

Bond Polarity• A similar diagram describing Bond Polarity is:

Electronegativity Difference Bond Type

Electronegativity (Again!)The Electronegativity of an atom

The relative attraction of that atom for the shared electrons in a bond.

Higher electronegativity means greater electron attraction to that atom

Another Schematic of the Pauling Electronegativity Scale

Schematic Relationship BetweenElectronegativity & Bond Type

Mixed Covalent & Ionic Bonding• Based on the electronegativity differences of the atoms in a bond,

some bonds have more ionic character (complete charge separation) than others. So, some bonds have some ionic character, but also have some covalent character (sharing electrons). The figure plots the ionic character of some bonds as a function of the difference in electronegativity of the two atoms involved.

CovalentBonds

IonicBonds

Table 3.8:Fractional Ionic Character of Bonds in Binary Crystals

The “Continuum” of Bonding TypesCovalent Ionic

Equal e- Sharing: Pure Covalent

• l

“Slightly” e- Unequal Sharing

• l

“Very” Unequal e- Sharing

Electron Transfer: Pure Ionic+ Ion - Ion

Metallıc Bondıng• Metallic Bonding is found in solid metallic elements. It is

caused by the electrostatic force of attraction between positively charged ions & delocalized valence electrons.

Metallic Bonds:Are typically weakerthan either ionic or

covalent bonds.

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In metals, the valence electrons are relatively weakly bound to the nuclei. So, they can move “freely” through the metal & they are spread out among the atoms in the form of a low-density “electron cloud”. A metallic bond results from the sharing of a variable number of electrons by a variable number of atoms. A metal may be crudely described as a cloud of “free electrons” moving in a lattice of positive ions. Therefore, metals have high electrical and thermal conductivity.

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+

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+

+

+

+

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Metallic Lattices are typically relatively empty.• That is, there are large internuclear spacings. The preferred

lattice arrangements are such that each atom has as many nearest neighbors as possible. The weakness of the individual bonding in a metal is due to this enlargement of the internuclear spacing.

The Valence Electrons in a Metalcombine to form a “sea” of electrons

that move relatively freely between the atom cores. The more electrons there are, the stronger the attraction.

• This means melting & boiling points are higher than for non-metals, & the metal is stronger & harder than non-metals. 

The positively charged cores are held together by these negatively charged electrons.

• The “free” electrons act as the bond (“glue”) between the positively charged ions. This type of bonding is nondirectional and is rather insensitive to structure.

• So, metals have a high ductility - the “bonds” do not “break” when atoms are rearranged – so metals can experience a significant degree of plastic deformation.

Hydrogen Bondıng• A H atom has only 1 electron, so it can be covalently

bonded to only one other atom. However, the H atom can involve itself in an additional electrostatic bond with a second H atom of highly electronegative character such as F or O. This 2nd bond permits a

HYDROGEN BONDbetween two atoms or structures.

• The hydrogen bond strength varies from 0.1 to 0.5 eV/atom.

• As already mentioned, H bonds connect water molecules in ordinary ice. H bonding is also very important in proteins & nucleic acids & therefore in life processes. As also already mentioned,

H bonding is a special case ofVan der Waals bonding.

Bonding Types

Ionic Bonding

Van Der Waals Bonding

Metallic Bonding

Covalent Bonding

Hydrogen Bonding

High Melting

Points

Hard & Brittle

Non-Conducting

NaCl, CsCl

ZnS

Low Melting

Points

Soft & Brittle

Non-Conducting

Ne, Ar, Kr, Xe

Variable Melting

Points

Variable

Hardness

Conducting

Fe, Cu, Ag

Very High

Melting Points

Very Hard

Usually not

Conducting

Diamond,

Graphite

Low Melting

Points

Soft & Brittle

Usually

Non-Conducting

Ice,

Organic Solids