Chemical Bonding Why do bonds form? to lower the potential energy between positive and negative charges positive charges protons cations negative charges electrons anions non-metals metals gain e- lose e- Periodic Table

Lewis electron-dot symbols element symbol = nucleus + core e- one “dot” = valence e- metals dot = e- it loses to form cation non-metal unpaired dot = e- paired through e- gain or sharing

Ionic bonding metal + non-metal Na (s) 2 + Cl2 (g)  NaCl (s) : . + Cl- : Na + Cl : . Na+ : [Ne] 3s1 [Ne] 3s23p5 [Ne] [Ar] 2 Ca(s) + O2(g)  2 CaO (s) . . . Ca : O : . [Ar]4s2 [He]2s2 2p4 Ca2+ O2- [Ar] [Ne]

high Electron Affinity Ionic bonding metal + non-metal low Ionization Energy high Electron Affinity lose 1 or 2 valence e- gain e- electron transfer takes place electrostatic attraction between cation and anion e- - + formula = ratio of anions to cations

Ionic sizes isoelectronic series same # electrons 46 e- + isoelectronic series same # electrons 46 e- +49 +50 +51 ions get smaller

Ionic bonding metal + non-metal 2 Na (s) + Cl2 (g)  2 NaCl (s)

Ionic bonding metal + non-metal 2 Na (s) + Cl2 (g)  2 NaCl (s) exothermic heat given off negative Ionization Energy Na Na+ + 496 kJ/mol Electron Affinity Cl Cl- -349 kJ/mol Lattice Energy E = k Q1 Q2 -787 kJ/mol d -640 kJ/mol Coulomb’s law Na+ Cl- NaCl +

Ionic solids cation + anion + - metal non-metal sodium + chlorine NaCl chloride Na Cl + - 801o C lithium + oxygen Li O lithium oxide Li O 2 + 2- > 1700oC magnesium + nitrogen Mg N magnesium nitride Mg N 2+ 3- 3 2 strong interactions (ion-ion) high melting points

Transition metals more than 1 form except Ag+ Zn2+ Cd2+ Al3+ aluminum manganese + oxygen + sulfur Al Mn1+ Mn2+ Mn3+ Mn4+ S 3+ 2- Al S 2 3 aluminum sulfide Mn3+ O 2- Mn2O3 Mn4+ O2- MnO2 manganese(IV) oxide manganese(III) oxide

Covalent bonding non-metal + non-metal electrons shared between atoms high Ionization Energies high Electron Affinities electron density between the atoms distance between atoms = bond length formula = actual # atoms

Lewis structure : . : F : . : F . F : F + : [He]2s22p5 [Ne] e- not used in bonding lone pairs shared e- bonding pair shared equally between F : . . . H O : : O . H 1s1 [He]2s22p4 [He] [Ne] oxygen 2 lone pairs bonding pair not shared equally

Covalent bonding non-metal + non-metal H . H . H . + 1s1 1s1 [He]

share valence electrons = chemical bonds carbon + chlorine C Cl 4+ - Covalent compounds non-metal + non-metal share valence electrons = chemical bonds carbon + chlorine mono di tri tetra penta hexa hepta octa C Cl 4+ - CCl4 carbon tetra chloride nitrogen + oxygen N ? O 2- NO nitrogen monoxide 2+ NO2 nitrogen dioxide 4+ N2O4 dinitrogen tetroxide 4+

Covalent compounds N H NH3 nitrogen trihydride ammonia 3+ 1- H O H2O dihydrogen monoxide water 1+ 2- weak forces low m.p. 0.0oC Table 2.5 p. 62 Polyatomic ions NH4+ ammonium ClO3- chlorate OH- hydroxide MnO4- permanganate NO3- nitrate CrO42- chromate SO42- sulfate CO32- carbonate PO43- phosphate

Metallic Bonding metal + metal metals valence e- well shielded low Ionization Energy low Electron Affinities share valence e- not localized between atoms delocalized move freely throughout metal Na (nucleus and core e-) e- “sea” (valence e-)

Electronegativity ability of an atom in a molecule to attract e- to itself related to Ionization Energy Electron Affinity Pauling scale

NaCl 2.1 ionic 801oC BeCl2 1.5 polar covalent 405oC AlCl3 1.5 PCl3 0.9 polar covalent 76oC Cl2 0.0 covalent non-polar covalent  0.4 C H polar covalent 0.5-1.8 + C O - ionic > 1.8 Li2O