chemical bond - attraction between atoms or ions

• ionic bond - electrostatic forces existing between ions of opposite charge
  • ions formed when electrons transferred
  • ionic compounds made from metals and nonmetals
• covalent bond - sharing of electrons
  • interactions between nonmetallic elements
• metallic bond - attraction between metals
  • each atom bonded to many neighboring atoms
  • bonding electrons free to move throughout the substance

ionic bonding - ions held in 3D array

• in forming ionic compounds, 1 atom loses an electron while another gains 1
• formation of ionic compounds usually exothermic
  • ionic compounds stable due to attraction of opposite charges
  • energy released by ion attraction makes up for endothermic ionization
• lattice energy - energy required to completely separate a mole of ionic compound into its ions
  • increases as charges on ions increase, radii decrease
• makes ionic compounds very hard/brittle, w/ high melting points
• covalent bonds hold atoms in polyatomic ions together, but polyatomic ions as a whole still act as ions
• E = k(Q1Q2)/d
  • Q1, Q2 = charges on the particles
  • d = distance between centers
  • k = constant 8.99 x 109 J-m/C2
  • energy increases as charge increases, or as distance between centers decrease

electron configurations of ions of representative elements

• by octet rule, ions tend to have electron configurations of noble gases
• electrons lost from subshell w/ highest n value first

electron configurations of metal ions

• usually only have net charges of +1, +2, or +3
• tries to have a full d subshell, loses electrons in s subshell first (higher n value)

covalent bonding - sharing of electron pairs

• nuclei’s attraction to the shared electrons overcomes their repulsion to each other
• single line used in Lewis structures to represent each shared electron pair
  • single bond - only 1 pair of electrons shared
  • double bond - 2 pairs of electrons shared, represented by 2 lines
  • triple bond - 3 pairs of electrons shared
  • multiple bonds shorter/stronger than single bonds
• # of shared electron pairs increase >> distance between bonded atoms decrease
• metals w/ high oxidation numbers tend to act molecular instead of ionic