Covalent Bonding

The Covalent Bond Octet Rule Covalent Bonds molecule The chemical bond that results from the sharing of valence electrons. Shared electrons count for both atoms Mostly in nonmetals molecule

Naturally occur as diatomic molecules Hydrogen - H2 Nitrogen – N2 Oxygen – O2 Fluorine – F2 Chlorine – Cl2 Bromine – Br2 Iodine – I2

Nature of Bonding Lewis Structures Single covalent bonds Use electron-dot diagrams to show how electrons are arranged in molecules Single covalent bonds 1 pair of shared electrons Double covalent bonds 2 pair of shared electrons Triple covalent bonds 3 pair of shared electrons

Sigma Bonds (s) Pi bonds (p) Electrons shared in between atoms Single covalent bond One in every molecule Pi bonds (p) Parallel orbitals overlap to share electrons In multiple bonds all but one

Structural formulas (Lewis Dot Structures) Predict the location of certain atoms. Hydrogen is always terminal. Least electronegative is the central atom Find the total # of valence electrons Determine the # of bonding pairs Place one bonding pair between the central atom and all of the terminal atoms Determine the # of remaining electron pairs. Place lone pairs around terminal atoms to satisfy the octet rule. Remaining pairs go to the central atom. Convert to multiple bonds in order to fulfill the octet rule if necessary.

Naming Binary Molecular Compounds The first name in the formula is always named first, using the element name from the periodic table. The second element in the formula is named using the root of the element and adding the suffix –ide. Prefixes are used to indicate the number of atoms of each type present in the compound. Mono- is not used on the first element.

Naming Binary Acids Use the prefix hydro- for the hydrogen part Use the root of the second part of the acid and add the ending –ic. End by adding the word acid to the end.

Naming Oxyacids Do not write anything for the hydrogen. Use the root of the oxyanion If the oxyanion ends in –ate replace it with –ic. If it ends in –ite replace it with –ous.

Laws Relating to Naming Law of definite proportions For any chemical compound, the masses of the elements are always in the same proportions. Always simple whole number ratios Law of multiple proportions Different masses of one element combine with the same masses of another element in a ratio of small whole numbers Form different compounds

Bond theories Molecular orbitals VSEPR theory Overlapped atomic orbitals in two atoms that are combined Orbitals apply to the molecule as a whole VSEPR theory Repulsion between pairs of electrons adjust to provide the greatest distance between electron pairs

Hybridization Atomic orbitals mix the result is hybrid orbitals

Polarity Sharing is not necessarily equal Table 8.3 page 238 Equal is nonpolar Unequal is polar Slight charge variations Table 8.3 page 238 Polar Molecules Two poles - dipole

Intermolecular Forces Weaker than Ionic or Covalent Determine State of molecular compound Van der Waals forces Two weakest intermolecular Dispersion Forces Weakest of all Moving electrons Greater in with more electrons

Hydrogen Bonds Dipole Interactions Strongest intermolecular Polar molecules Charges in dipoles Hydrogen Bonds Strongest intermolecular 5% strength of a covalent bond Polar bonds with hydrogen