Chemical Bonding

• Chemical Bond


  A bond results from the attraction of nuclei for electrons
• Atoms


  • All atoms are trying to achieve a stable octet
• Chemical bonds


  Electrostatic forces of attraction that holds atoms together
• Octet


  Stable with eight electrons
• Duet


  Stable with two electrons (i.e. Helium and Hydrogen)
• Isoelectronic


  Refers to two atoms, ions, or molecules that have the same electronic structure and the same number of valence electrons
• Isoelectronic


  • Ca+² is isoelectronic with Argon
  • O-² is isoelectronic with Neon
• Three Major Types of Chemical Bonding

  • Ionic Bonding
  • Covalent Bonding
  • Metallic Bonding
• Valence electrons


  The electrons on the outermost shell of an atom, responsible for the formation of a chemical bond
• Lewis symbol


  Depicts the valence electrons of an atom as dots around the symbol for the element
• The number of valence electrons determines the group placement of an element
• Valence electrons


  • Hydrogen has one valence electron, placing it in the alkali metal family
  • Fluorine has seven valence electrons, placing it in the halogen family
• Ionic Bonding


  Always formed between metal cations and non-metals anions, where oppositely charged ions stick like magnets
• Ionic compounds


  Neutral compounds made up of positively charged ions (cations) and negatively charged ions (anions)
• How ionic compounds form

  Valence (outermost) electrons of one atom are transferred permanently to another atom, with the atom that loses electrons becoming a positively charged ion (cation) and the one that gains them becoming a negatively charged ion (anion)
• Ionic compounds


  • Sodium must give up one valence electron to resemble the electronic configuration of neon, its nearest noble gas
  • Chlorine must gain one electron to attain a stable configuration resembling argon
• Covalent Bonding


  The stable balance of attractive and repulsive forces between atoms when they share electron pairs
• Covalently bonded atoms


  Form MOLECULES
• Electron-Dot Diagrams


  Combine the valence electrons of atoms to show the covalent bonds
• General Rules for Drawing Lewis Structures


  • All valence electrons of the atoms must be shown
  • Generally each atom needs eight electrons in its valence shell (except Hydrogen needs only two electrons and Boron needs only 6)
  • Multiple bonds (double and triple bonds) can be formed by C, N, O, P, and S
  • Central atoms have the most unpaired electrons
  • Terminal atoms have the fewest unpaired electrons
  • Lone pair - unpaired electrons
  • When carbon is one of the atoms, it will always be in the center
  • Sometimes there is no central atom

  
• Steps to determine the number of bonds formed between atoms

  1. Get the total valence electrons in a compound

  2. Compute the Octet Requirement that each atom should have 8 valence electrons to become stable

  3. Subtract the total available valence electrons from the octet requirement, then divide the difference by 2 because a pair of shared electrons is equal to 1 bond
• Polar and Non-polar Covalent Bonds

  Polar covalent bonds have uneven sharing of electrons, while non-polar covalent bonds have equal sharing of electrons
• Electronegativity


  The ability of atoms to attract electrons in a chemical bond
• Electronegativity values


  • H = 2.1, B = 2.0, C = 2.5, N = 3.0, O = 3.5, F = 4.0 (highest), Cs = 0.7 (lowest)
• Determining bond types using electronegativity

  Difference in electronegativity ≥ 0.5 = Polar, Difference in electronegativity ≤ 0.4 = Non-polar

  
• Direction of Polarity


  Electron will go to the element with higher electronegativity
• Dipole


  Molecule or covalent bond that has a separation of charges