Chapter 9: Ionic and covalent bonding

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Chapter 9 Ionic and covalent bonding
Chemistry 1061 Principles of Chemistry I Andy
Aspaas, Instructor
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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 (Ne3s1) Cl (Ne3s23p5) ? Na (Ne)
  Cl (Ne3s23p6)
• 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

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Signifying ionic bond formation

• Lewis electron-dot symbols valence electrons
  (electrons in outer shell) represented by dots
  drawn around atoms 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

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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 Coulombs 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)

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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

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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

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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)

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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

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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

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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)

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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

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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