Chemical Bonding

Chemical Bonding

1. Intermolecular Forces

   • Electrostatic Forces: The forces present in ionic bonds, where one atom is positively charged, and the other is negatively charged.
   • Dipole-Dipole forces occur with two polar molecules, and the partial charges that come with them. Dipole-Dipole forces are very weak, and this is how molecules bond with each other. These forces are also permanent.
     • Hydrogen Bonding: A weak force that forms a special Dipole-Dipole bond. This occurs when a hydrogen atom connected to an electronegative atom is in the vicinity of another electronegative atom with a lone pair of electrons. Stronger than Dipole-Dipole forces, but weaker than ionic bonds. It isn’t really a bond
     • LDF’s or London Dispersion Forces are the forces between electrons and the random movement of electrons. This random movement created temporary Dipoles, creating the weakest attraction and charge. The more electrons an atom or molecule has, the stronger the LDF’s.

2. Electronegativity and Bonding

   • Electronegativity is the measure of an atom’s tendency to attract an electron. The difference in electronegativity between two elements determines the bond that forms between them.
     • Ionic bonding is the complete transfer of electrons from one atom to another, and occurs when the EN difference is above 1.7. When Ionic bonds form, two oppositely charged ions appear.
     • The Pauling Chart, the most common chart used to determine electronegativity
     • Covalent bonding happens when the difference in electronegativity is low enough that the elements share the electron equally. This usually happens when the EN difference is less than or equal to 0.4.
       • Polar Covalent bonding happens when there is unequal sharing of electrons between elements. The bond isn’t completely ionic, but there is a partial charge on one side of the molecule. This happens when the EN difference is between 0.4 and 1.7.
       • Hybridization is when valence electrons in various atoms rearrange themselves in order to form more covalent and polar covalent bonds. Electrons will move up into separate subshells and orbitals in order to make sure that the element has a completed octet.

3. Diagrams

   • Lewis Dot Diagrams: A drawing of a molecule that shows all atoms involved, as well as the lone pairs, and any bonds that may form between them.
     • Lewis Structure is how molecules are drawn and represented. Lines are used to represent covalent bonds, and each atom needs a completed octet
   • Molecular Shape is how molecules will arrange themselves in order to ensure that valence electron pairs are as far apart as possible.
     • A coordinate covalent bond is when one atom provides both of the electrons in a covalent bond so that both atoms reach a complete octet. An example of this is CO, where oxygen provides one pair of electrons as well as the electrons already shared.
       • VSEPR stands for valence-shell-electron-pair-repulsion theory. This is the theory for how molecules arrange themselves. This theory sums up molecular shape and the rules for said shape.
         • Molecular Shape is how molecules will arrange themselves in order to ensure that valence electron pairs are as far apart as possible.
           • Molecular polarity is the electronegativity of all the bonds in a molecule. A molecule will be polar if there is a greater mass of electrons in one part of the molecule.