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

2. 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.

3. Covalent Bonding of 2 H Atoms  The H 2 Molecule Interaction Potential

4. 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

5. 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. Brittleness They are hard because the atoms are strongly bound in the lattice, and are not easily displaced. Hardness Poor conductors because electrons are held either on the atoms or within covalent bonds. They cannot move through the lattice. Electrical conductivity 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 point ExplanationProperty Covalent Materials

6. Pictorial Comparison of Ionic & Covalent Bonding

7. In a Covalent Bond, neighboring atoms SHARE 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 electronic configuration for each atom. Covalent Bonds

8. In an Ionic Bond, neighboring atoms TRANSFER electrons from one site to another. Positive & negative ions are formed. In solids, an ionic solid or a salt is formed. Example: NaCl Na o  Na + (cation) + e - Cl o + e  Cl - (anion) Na o + Cl o  NaCl (salt) Ionic Bonds (Contrast to Covalent Bonds)

9. Mixed Covalent & Ionic Bonding NOTE!! Some bonding in many materials Cannot 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.

10. 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

11. Bond Polarity A similar diagram describing Bond Polarity is: Electronegativity Difference Bond Type

12. 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

13. Schematic Relationship Between Electronegativity & Bond Type

14. 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. Covalent Bonds Ionic Bonds

15. Table 3.8: Fractional Ionic Character of Bonds in Binary Crystals

16. The “Continuum” of Bonding Types Covalent Ionic Equal e - Sharing: Pure Covalent l “Slightly” e - Unequal Sharing l “Very” Unequal e - Sharing Electron Transfer: Pure Ionic + Ion - Ion

17. 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 weaker than either ionic or covalent bonds.

18. v 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. + + + + + + + + + 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.

19. The Valence Electrons in a Metal combine 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.

20. 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 2 nd bond permits a HYDROGEN BOND between two atoms or structures. The hydrogen bond strength varies from 0.1 to 0.5 eV/atom.

21. 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 of Van der Waals bonding.

22. Bonding Types Ionic Bonding High Melting Points Hard & Brittle Non-Conducting NaCl, CsCl ZnS Van Der Waals Bonding Low Melting Points Soft & Brittle Non-Conducting Ne, Ar, Kr, Xe Metallic Bonding Variable Melting Points Variable Hardness Conducting Fe, Cu, Ag Covalent Bonding Very High Melting Points Very Hard Usually not Conducting Diamond, Graphite Hydrogen Bonding Low Melting Points Soft & Brittle Usually Non-Conducting Ice, Organic Solids