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Bonding: General Concepts

Bonding: General Concepts. Chapter 8. Words to know:. chemical bond ionic bond covalent bond metallic bond Lewis Symbol Octet rule. Practice. Classify the following compounds as ionic or covalent. Justify your answer. Which compounds contain both types of bonds? - KBr - SO 2

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Bonding: General Concepts

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  1. Bonding: General Concepts Chapter 8

  2. Words to know: • chemical bond • ionic bond • covalent bond • metallic bond • Lewis Symbol • Octet rule

  3. Practice Classify the following compounds as ionic or covalent. Justify your answer. Which compounds contain both types of bonds? - KBr - SO2 - H2SO4 - CH3COOH - Na3PO4 - CaCO3

  4. Lewis Symbol • Element symbol with valence electrons written around it as dots • Elements want to gain, lose, or share electrons to look like a noble gas (isoelectronic)

  5. Place the following chemical species into isoelectronic groups: N3-, K+, Ca2+, O2-, F-, Ne, Br, Kr, Sc3+, Na+, Al3+, Se2-, Mg2+

  6. Ionic Bonding • electrons are transferred from an atom with low electronegativity to one with high EN • electrostatic attraction between the two oppositely charged ions • arranged in a crystal lattice (lattice energy, DHL, is energy required to completely separate solid ionic compound into gaseous ions)

  7. Lattice Energies? • Lattice energy increases when the atoms are smaller or have a higher charge (exchanging more electrons) • higher lattice energy means the ionic compound is more strongly bonded • high lattice energies also explains why ionic compounds are brittle and hard

  8. In each of the following pairs of compounds, identify the one with the higher lattice energy • KCl, CaS • LiF, NaCl • Fe2O3, MnO2 • CaO, CaCl2

  9. Ions • Representative elements follow the “hill of oxidation numbers” when in ionic compounds • transition metals (including lead & tin) are a little weird: • Their valence electrons are the highest filled s sublevel and occasionally 1 or more of their d electrons

  10. Transition Metal Ions • Iron • Fe: [Ar] 4s23d6 • Fe2+: [Ar] 3d6 • Fe3+: [Ar] 3d5 • Lead • Pb: [Xe] 6s24f145d106p2 • Pb2+: [Xe] 6s24f145d10 • Pb4+: [Xe] 4f145d10

  11. Write the electron configurations for Cr3+ and Sn4+

  12. Covalent Bonding • 2+ atoms sharing electrons • Lewis structures show shared and lone pairs of electrons • polar or non-polar determined by difference in EN values for the 2 bonding elements • 0-0.4 = NPC • .4-1.0 = PC • 1.0-2.0 = really PC • >2.0 = ionic

  13. Dipoles • With polar covalent bonds, there is a dipole (one end of the bond hogs the “shared” electrons a little more than the other) • symbolized with d+ and d- and are not whole number charges

  14. Drawing Lewis Structures • Add up the total valence electrons of all bonding atoms • Use one pair of electrons to bond each outer atom to the central atom (usually the least EN or the one present in the least abundance) • Complete octets around all of the outer atoms • Place any remaining electrons around the central atom • If there aren’t enough electrons to give the central atom an octet, make multiple bonds

  15. Practice Write the Lewis structures for each of the following compounds: • NO3- • CO2 • PCl5 • NO3

  16. Resonance Structures • used when 2+ Lewis structures are equally good representations of the bonds • actual structure is kind of an average of all the possibilities • examples include ozone & nitrate ion

  17. Exceptions to Octet Rule • When there is an odd # of electrons, one atom will only have 7 electrons around it. (NO, NO2) • When the compound has a group 2 or 3 element as the central atom, the number of electrons around the central atom will be twice the group number • If central element is big, it can have an expanded octet (PCl5)

  18. Covalent Bond Strength • Multiple covalent bonds are stronger than single covalent bonds (they are also shorter) • Higher the number of electrons shared, the stronger the bond

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