Periodic Properties of the Atom Properties that depend on position of element in the periodic table. Factors that affect the periodic properties: 1. Principal quantum number of valence shell (n valence ) * within a column/group...as proceed down column n valence increases and shell size increases. * within a row/period..as proceed across row n valence remains constant. 2. Effective nuclear charge (Z eff ): reduced nuclear charge felt by valence e- in an atom due to shielding by core e-. * within column/group..as proceed down column Z eff remains relatively constant. * within a row/period..as proceed across a row Z eff increases. Z eff = Z - # core e- Compare the effective nuclear charge felt by the valence electrons in sodium vs. chlorine. Periodic Properties A. Atomic Radius: distance from nucleus to outer limits of atom. Note: Electron distribution on average is considered to be spherically symmetric about nucleus. * Why do atomic radii/sizes increase as proceed down periodic table? Why do atomic radii/sizes decrease as proceed across periodic table? Compare the radii of F, C, N, Ne, and O. B. Ionic Radii 1. Cation < Neutral Atom 3. Isoelectronic species Which is larger, Na + or Na? Explain. Relative radii depend on Z eff. Which is larger, F -, Ne, or Na +? 2. Anion > Neutral Atom Which is larger, F - or F? Explain 1
C. Ionization Energy (IE): Energy absorbed when electron is removed from the valence shell of gaseous atom/ion in ground state. Note: Harder is to remove e-, the larger the IE. In general.. nonmetals, have higher IE relative to metals, which have lower IE. Example: Consider oxygen. O(g) O + (g) + 1 e- IE 1 (First IE) O + (g) O +2 (g) + 1 e- IE 2 (Second IE) Compare the IE of Li, F, and N. Compare the IE of Ar, He, and Ne. D. Electron Affinity (EA): Energy released when an electron is added to the valence shell of a gaseous atom/ion in its ground state. Note: Easier is to add e-, the more energy released (i.e. more exothermic). In general.. nonmetals, have most negative EA relative to metals, which have less negative EA. Example: Consider oxygen. 1 e- + O(g) O - (g) EA 1 =- 1 e- + O - (g) O -2 (g) EA 2 =+ Compare the EA of Mg, Cl, and Si. Compare the EA of K, Na, and Li. Of special note: Ionic Compound = Metal + Nonmetal (Low IE) (Most neg. EA) easy to remove e- easy to add e- to form cation to form anion 2
Electronegativity and Ionic vs. Molecular Compounds Ionic Compounds: formed by combination of metal cation with nonmetal anion. Ions held together by ionic bonds (due to electrostatic attraction between ions of opposite charge). Molecular Compounds: Formed by combination of two or more nonmetals. Atoms held together by covalent bonds (due to sharing of pair of e- between atoms). Problem..actual dividing line between ionic and covalent compounds is not so clear cut!! Many compounds have both ionic and molecular properties. Linus Pauling (American scientist, Nobel prize in Chemistry 1954) Studied the nature of chemical bond. Defined concept of electronegativity (en): ability of bonded atom to attract e- in chemical bond toward itself. Note: The higher the en, the more the bonded atom pulls e- toward itself. In general..metals have lower en, nonmetals have higher en. Exception, Group VIIIA, en=0. Electronegativity increases en=0 for Group VIIIA Electronegativity increases Use of Electronegativity to Determine Bond Type 1. Look at difference in electronegativity, en = en 2 en 1. 2. If en 2, then ionic bond If en = 0, then covalent bond. If 0 < en < 2, then polar covalent bond (bond intermediate between ionic and covalent and with partial charge separation). 3
* What type of bonds are present in each of the compounds shown below? (en: Ca=1.1, O=3.5, C=2.5, H=2.1, Cl=2.9, S=2.4, Zn=1.7) * If polar covalent bonds are present, indicate which atom has the partial negative and which the partial positive charge. CaO CH 4 Cl 2 SO 4-2 ZnS Lewis Dot Symbols Used to keep track of valence electrons. To draw Lewis Dot Symbols: 1. Write chemical symbol for element. 2. Surround by dots, each dot corresponding to a valence e-. Remember..for A-Group elements, # valence e- = Group #. A. Draw Lewis Dot Symbols for Mg, S, Cl, and Ar. B. Use Lewis Dot Symbols to diagram the reaction between Na and F to form the ionic compound NaF. C. In NaF, do the elements obey the octet rule? State the octet rule. Which elements tend to obey the octet rule? Which elements are exceptions to the octet rule? Do atoms in molecular compounds also tend to obey the octet rule? 4
Lewis Structures Use Lewis Dot Symbols to show the arrangement of valence e- such that octet rule is obeyed. A. Draw Lewis structure for F 2, O 2, and N 2. B. Define the following terms: single bond, double bond, triple bond, bonding e- pair, nonbonding (or lone) e- pair, and bond order. Lewis Structures: Steps for Drawing 1. Count up total # valence e-. 2. Determine skeletal structure (what s bonded to what). a. Central atom is usually written first in chemical formula. b. H is never the central atom. 3. Place 2 e- in each bond. 4. Surround non-central atoms by octet of e-. a. H is exception and is only surrounded by 2 e-. 5. Count up how many e- remain. Place any remaining e- on central atom as e- pair. 6. If central atom does not have octet of e-, form multiple bond(s) to it to give it an octet of e-. a. H, B, Be, F and metals do not generally form multiple bonds. 7. If central atom in Step #6 does not form multiple bonds, leave it short of e-. 5
A. Draw Lewis structures for the substances given below. CF 4 CH 4 CO CO 2 NCl 3 OCl 2 H 3 O + NO 3 - CO 3-2 ClO 2 - ClO 2 + HCN BeCl 2 BCl 3 PCl 5 SF 4 HClO 4 HNO 2 B. Why can S and P have more than 8 e- in the valence shell? Can elements such as C, N or O (Period 2 elements) expand their octet? Explain. Resonance Must be invoked when one single Lewis structure does not adequately describe the bonding in a molecule or ion. Resonance Structures: 1. Same skeletal arrangement. 2. Different arrangement of e- pairs. Example: A. Draw a Lewis structure for carbonate. Does this one Lewis structure accurately describe the bonding in carbonate? The actual carbonate structure is a resonance hybrid of all reasonable Lewis structures. What is a resonance hybrid? B. Draw two reasonable resonance structures for sulfuric acid. Use formal charge considerations to predict which form is more stable (or preferred). 6
Formal Charge Gives information on how many e- are still associated with an atom relative to the number of valence e- the atom came in with. Formal Charge = # Valence e- - # bonds - total # nonbonding e- Rules for determining stability of resonance structures: 1. Structure with least amount of formal charge separation is more stable. 2. Structure with negative formal charge on more en element is more stable. 3. Structure with two positive or two negative formal charges on adjacent atoms is NOT stable. A. Draw resonance structures for N 2 O. Which structure is most stable? least stable? B. Draw resonance structures for HClO 3. Which structure is most stable? least stable? 7