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Chapter 9: Ionic and covalent bonding

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

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Chapter 9: Ionic and covalent bonding

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

  2. 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

  3. 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 —

  4. 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)

  5. 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

  6. 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 ns2configurations (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

  7. 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)

  8. 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

  9. 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

  10. 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)

  11. 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

  12. 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

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