Chapter 9: Ionic and covalent bonding Chemistry 1061: Principles of Chemistry I Andy Aspaas, Instructor

Electron configuration of ions Ionic bonds: formed by electrostatic attraction between oppositely-charged ions Ions are normally formed by adding or removing electrons from atoms to give them a noble-gas configuration Consider formation of sodium chloride –Na ([Ne]3s 1 ) + Cl ([Ne]3s 2 3p 5 )  Na + ([Ne]) + Cl — ([Ne]3s 2 3p 6 ) –The oppositely charged sodium cation and chloride anion now have noble-gas configurations, and become ionically bonded –NaCl crystal involves an orderly arrangement of Na + and Cl — ions

Signifying ionic bond formation Lewis electron-dot symbols: valence electrons (electrons in outer shell) represented by dots drawn around atom’s element symbol –First put one dot on each of 4 sides, then add 2nd dot to each side, until all valence electrons are drawn Na + Cl  Na + + Cl —

Energy involved in ionic bonding Ionization energy: energy required for an atom to lose an electron –Positive value, but small for groups IA - IIIA Electron affinity: energy released when an atom gains an electron –Negative value, especially favorable for groups VIA - VII7A Ion pair energy: energy released when oppositely charged ions are brought into a pair (calculated by Coulomb’s law) Lattice energy: energy required to break a lattice of ions into gas-phase atoms (reverse is the energy released when forming gas-phase ions into a lattice)

Properties of ionic substances Ionic substances: normally high-melting solids –Due to strong attractions between ions which must be broken if the solid is to melt MgO has much higher melting point than NaCl, since each ion has 2+/2— charge instead of just 1+/1— Molten ionic substances conduct electricity, just like a solution with dissolved ions would

Predicting ion charges Groups IA & IIA form cations to give noble-gas configurations (charge = group #) Metals in groups IIIA - VA can form cations either with noble- gas configuration, or with ns 2 configurations (charge = group # or group # — 2) Nonmetals in groups VA - VIIA form anions with noble-gas configurations (charge = 8 — group #) Many transition metals form +2 charges by losing their two highest s electrons –+3 is formed by losing the two highest s electrons and one d electron

Covalent bonds Covalent bonds involve sharing of a pair electrons between two atoms –Ex.: Formation of H 2 : H· + ·H  H : H Electron pairs in Lewis electron-dot formula can be either bonding pair (shared between two atoms) or a nonbonding pair (unshared, remains on one atom) Covalent bonds usually exist between nonmetals, where formation of an ion-pair would be unfavorable Octet rule: many atoms prefer 8 valence electrons available when forming covalent bonds (some do not)

Polar covalent bonds Electronegativity: ability of an atom in a molecule to draw bonding electrons to itself Fluorine is the most electronegative element, Cesium is the least –Electronegativity decreases as you go left, or down on the periodic table Uneven electronegativities of atoms involved in a covalent bond will yield uneven sharing of the electrons; this is a polar bond

Lewis structures Lewis structure: electron dot structure for an entire molecule Use dots to indicate unshared electrons and lines to indicate covalent bonds One line represents a single bond (2 shared electrons) –2 lines for a double bond, 3 for a triple bond, etc

Drawing Lewis structures Predict skeleton structure (atom arrangement) by choosing a central atom (usually least electronegative) Find the total number of valence electrons in the molecule (for a polyatomic ion, add an electron for a 1– charge, remove an electron for a 1+ charge) Count the bonds you have already drawn as pairs of valence electrons, and distribute remaining valence electrons as pairs among the surrounding atoms to satisfy octet rules Add remaining electrons as pairs to central atom If octets cannot be filled, try adding double or triple bonds (C, N, O, and S often form multiple bonds)

Exceptions to the octet rule Nonmetals in row 3 and beyond can use higher- energy empty d orbitals for bonding –Ex. PF 5 and SF 6 Also group IIA and IIIA atoms can form covalent compounds with less than 8 electrons in their valence shells –Ex. BF 3, BeF 2

Formal charges Hypothetical charges on individual atoms in a molecule Formal charge = valence electrons on free atom – 1/2 # shared electrons – # unshared (lone-pair) electrons If several Lewis structures are possible, the most important Lewis structure is the one with the fewest formal charges If two Lewis structures have the same number (and magnitude) of formal charges, choose the one with the negative formal charge on the more electronegative atom