Periodic Table - e • Columns: Similar Valence Structure - - - e e e - e inert gases inert give give up 1 accept 2 accept 1 H give up 2 He Li Be O F Ne Adapted from Na Mg give up 3 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 electrons Readily acquire valence electrons to become (+)ve ions. to become (-)ve ions. 1 Electronegativity Electronegativity (EN) •Tendency of an atom to attract/acquire an electron, i.e., accept valence electrons to become (-)ve ions (ionic bonding) or sometimes they share electrons with other atoms (metallic bonding). •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). •Atomic radius increases in size going down groups because new shell is added. •Atomic radius decreases from left to right because nucleus contains more protons (↑ Z)– addition of one proton 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. •Ionic radius: Cations (+) : smaller than atoms Anions (-) : larger than atoms •Ions always shrink with increasing positive charge and expand with increasing negative charge. •Mass increases with atomic number (increasing # of protons in the nucleus) 3 Summary of Periodic Trends A third class of elements, the metalloids, which straddle the metal–nonmetal 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 metals or nonmetals: 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 solids. •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 solid 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 in Solids (Bonding Forces and Energies) •Many physical properties can be predicted based on interatomic forces that bind the atoms together. •Net force 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 bond length. ~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 energy 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 separationbond 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 valence electron sharing between two adjacent atoms such that each atom assumes a stable electron configuration. Example is methane: 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 carbon sp3 valence electrons is shared for total of 8 valence electrons (Ne electron configuration), octet rule. •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 crystals 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(diamond) 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 Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University. • Molecules with nonmetals, e.g. Cl2, F2, O2 • Elemental solids (RHS of Periodic Table), 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 molecular solid 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 molecule are linked by covalent bonds, but the molecules that make up the crystal are held together only by the weak interactions known collectively as intermolecular forces or secondary bonds (including van der Waals, dipolar, and hydrogen 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 halogen, 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 charge density accumulates between the relatively positive atomic cores. •Defining characteristic of a covalent bond is the existence of a local maximum in the valence electron density 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 semiconductors, e.g.