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]3s1) + Cl ([Ne]3s23p5)

Na+ ([Ne]) + Cl— ([Ne]3s23p6)– 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 ns2

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 H2: 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. PF5 and SF6

• Also group IIA and IIIA atoms can form covalent compounds with less than 8 electrons in their valence shells

– Ex. BF3, BeF2

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