Periodic Table -

e • Columns: Similar Structure


- -


e e



inert inert

give give 1 up

accept 2 accept 1 accept H give 2 up He Li Be O F Ne Adapted from Na Mg give 3 up S Cl Ar Fig. 2.6, Callister & K Ca Sc Se Br Kr Rethwisch 3e. Rb Sr Y Te I Xe Cs Ba Po At Rn Fr Ra

Electropositive (more metallic) elements: Electronegative elements: Readily give up valence Readily acquire valence electrons to become (+)ve . to become (-)ve ions. 1

Electronegativity (EN) •Tendency of an to attract/acquire an , i.e., accept valence electrons to become (-)ve ions () or sometimes they share electrons with other (). •Thus, it’s opposite of electropositivity (metallicity) and its trends. •Atoms are more likely to accept electrons if their outer shells are almost full, and if they are less shielded from (closer to) the nucleus. •Numerous EN scales have been proposed, but the most commonly used is the one originally devised by Pauling. EN is on a 0 to 4.0 scale, with F having the highest EN of 4.0 and Cs and Fr having the lowest with 0.7.

Larger electronegativity (smaller electropositivity or metallicity) Smaller electronegativity (larger electropositivity or metallicity) 2 2 Size and Mass

Size and Mass •Trends are the same as in electropositivity (metallicity). • increases in size going down groups because new shell is added. •Atomic radius decreases from left to right because nucleus contains more (↑ Z)– addition of one to an atom gives a stronger nuclear (greater positive) charge to the nucleus and thus the electrons are attracted more tightly to the nucleus, thus atomic radii decreases.

: Cations (+) : smaller than atoms Anions (-) : larger than atoms •Ions always shrink with increasing positive charge and expand with increasing negative charge.

•Mass increases with (increasing # of protons in the nucleus)

3 Summary of

A third class of elements, the , which straddle the boundary, Zintl line (e.g.: B, Si, Ge, As, Te, and Se).

4 Bonding Generalizations Based on Periodic Trends in EN •We begin by classifying all elements as either or : the ‘bold’ stepped Zintl line across the right hand side of the previous slide (table). •With elements divided up in this fashion, we establish the following rules. •First, metallic elements and metal–metal combinations form metal bonded . •Second, nonmetallic elements (e.g. Cl-Cl), and nonmetal–nonmetal (e.g., O-N) combinations are generally covalently bonded. •Third, bonds between metals and nonmetals are either ionic or covalent, depending on the electronegativity difference which requires calculation: •This EN difference from Pauling’s expression for the ionicity fraction of a bond (f ) is 2 -1/4(xnm-xm) f=1-e exs.: binary ZnO & BaF2… where xnm is the EN of the nonmetallic element and xm is the EN of the metallic element. •We will assume that when f > 0.5, the bonds are ionic and that when f ≤ 0.5, the bonds are covalent. In generic terms, the greater the difference between two atoms, the more ionic the bond, while the smaller the difference the more covalent the bond. •In ternary or more complex compounds, the fractional ionicity can be determined by using stoichiometrically weighted averages for the values of xm and/or xnm, ex.: In4Sn3O12…

•It must be emphasized that the change from metallic to nonmetallic character is continuous and complex, so much that many authors would refute the apparently arbitrary binary categorization defined above…. 5 Bonding – Property Interrelationships

•However, with such criticism noted, a binary classification is nevertheless implemented because it has the practical advantage of leading to a simple set of rules to determine bond types. •Once the bond type is defined, the type of atomic structure and properties that the might have can also be inferred. This relationship is illustrated schematically:

•When crossing metal-nonmetal boundary, properties change dramatically with variation in bonding:

6 Origin of (Bonding and ) •Many physical properties can be predicted based on interatomic forces that bind the atoms together.

•Net FN between atoms = FA + FR which is also function of interatomic separation.

•Equilibrium exists when FA + FR become equal, there is no net force (=0), or else the atom moves. •The center of atoms will remain separated by equilibrium spacing (ro) or . ~0.3 nm •Potential energies between atoms

E=-∫Fdr or EN  EA  ER Adapted from Fig. 2.8(b), •The net E curve has a potential trough or Callister & Rethwisch 3e. well around its minimum. Here ro corresponds to separation distance at minimum of E curve.

•The bonding energy (E0) for two atoms corresponds to energy at this minimum point.

•E0 represents the energy required to separate these two atoms to an infinite separationbond strength

•The magnitude of E0 and shape of the curve vary 7 from material to material & depend on type of 7 atomic bonding. ….. Review: Covalent Bonding Model

•Covalent - there is sharing between two adjacent atoms such that each atom assumes a stable . Example is : C: has 4 valence e-, needs 4 more •Each atom contributes at least one electron to the bond and the H: has 1 valence e-, needs 1 more shared electrons may be considered to belong to both atoms. •For methane each H atom can acquire a He electron configuration, when one of four sp3 valence electrons is shared for total of 8 valence electrons (Ne electron configuration), . •The number of covalent bonds possible for a particular atom is determined by the number of valence electrons. For N' valence Adapted electrons, an atom can covalently bond with at most 8-N' other atoms. from Fig. •Examples: Cl atom has valence electron structure of 3s23p5 2.10, Callister & Rethwisch •C and Si have 2s22p2 and 3s23p2 structures, respectively. 3e. •f ≤ 0.5 (small DEN) metallic-nonmetallic elements and in with nonmetallic elements. •Bonds are directional in nature: they exist only in the direction between one atom and another that participate in electron sharing. •Thus covalently bonded materials are generally less dense than ionically or metallically bonded ones (non-directional); when bonds are directional, the atoms cannot pack together in as dense a manner, thus yielding a lower mass density. 8 Examples of Covalent Bonding

H2O H 2 F2 C() H He - 2.1 SiC column IVA Cl2 Li Be C O F Ne 1.0 1.5 2.5 2.0 4.0 - Na Mg Si Cl Ar 0.9 1.2 1.8 3.0 - K Ca Ti Cr Fe Ni Zn Ga Ge As Br Kr 0.8 1.0 1.5 1.6 1.8 1.8 1.8 1.6 1.8 2.0 2.8 - Rb Sr Sn I Xe 0.8 1.0 1.8 2.5 - Cs Ba Pb At Rn 0.7 0.9 1.8 2.2 - Fr Ra Adapted from Fig. 2.7, Callister 6e. (Fig. 2.7 is 0.7 0.9 GaAs adapted from , The Nature of the , 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University.

with nonmetals, e.g. Cl2, F2, O2 • Elemental solids (RHS of ), e.g. C, P, S, I • Compound solids with metals and nonmetals, e.g. ZnO, GaAs,GaP • Compound solids (about column IVA), e.g. SiC 9 Simple Bonding Models – Covalent (continued)

The has covalent bonds (dark lines) only within individual There is a covalently bonded molecules. Thus, there is no path between any two atoms, covalently bonded path e.g. 1 and 2. between the atom labeled 1 and the atom labeled 2; the molecules are bonded to one another only by weak Covalently bonded 3-D network such as Si, secondary forces. SiC, BN, etc. Molecular solids or polymeric solids such as •All atoms are linked by covalent bonds, i.e., crystalline materials C60, H2O, and there is a covalently bonded path between any macromolecular solids polyethylene. 2 atoms in the solid. •Atoms within each are linked by covalent bonds, but the molecules that make up the are held together only by the weak interactions known collectively as intermolecular forces or secondary bonds (including van der Waals, dipolar, and bond). •In such solids, not all atoms are connected by a path of strong covalent bonds. •“Rule of Thumb”: if more than two thirds of the components in a covalently bonded compound are H, C, O, N, or a , then it is likely to be a molecular solid. •However, diamond is a noteworthy example illustrating that this guideline should be applied 10 with caution. Covalent Bonding Model (continued)

•Simple model assumes that electrons are shared between atoms and that electron accumulates between the relatively positive atomic cores. •Defining characteristic of a is the existence of a local maximum in the valence in the regions between the atomic cores. For example, experimentally measured charged density in Si is shown: •The peak in electron density (.69) at the midpoint connecting the two Si nuclei is signature of the covalent bond. Fig. Valence electron density map in the •Concentrating the valence electrons in the spaces between {110} plane of Si. Contours are at 0.1 e/Å3. The shape of the peaks are the atomic cores is clearly distinct from the ionic bonding theoretically predicted and also found in model, where the valence electrons are centered on the anion many III-V , e.g. GaAs. positions, and the metallic bonding model where the valence electrons are uniformly distributed in the free electron “sea.”

11 Review: Metallic Bonding Model

•Metallic - Metallic materials have one, two or at most three valence electrons. •With this scheme, these electrons are not bound to any particular atom in the solid and are more or less free to drift throughout the entire metal sharing electrons.

•The remaining nonvalence electrons and atomic nuclei form cores, which possess a net positive charge equal in magnitude to total valence electron charge per atom.

•The ion cores arranged periodically are shielded from one another, and also "glued" together by the “sea” of valence (free) electrons or electron clouds. •In other words, the free electrons shield the positively charged ion cores from mutually repulsive electrostatic forces, which they would otherwise exert upon one other, thus metallic bond is non-directional.

•Due to large number of freely moving electrons, metals are good thermal (conduction of heat by free electrons) and electrical 12 conductors.