Review Sheet for Bonding Unit
Important Terms :
Bond
Rule of Octets
Rule of Duets
Stability
Isoelectronic
Polyatomic ion
Molecule
Electron
Valence Electron
Valence Shell
Lewis Dot structure
Odd numbered electrons
Structural Formula
Molecular Formula
Electronegativity
Pauling
Ionic bond
Polar covalent bond
Pure covalent bond
Co-ordinate covalent bond
Ionic character of a covalent bond
Bond angle
Bond length
Single bond
Double bond
Triple bond
Sigma bond (s)
Pi bond (p)
Central atom
Bonding pair (shared pair)
Lone pair
VSEPR Theory
Gillespie
Electron repulsions
"Electrons hate each other " --- 3 parts
General Form
Actual (molecular) Shape
Distortion
Bonding zones
Ligands
Bonding zones
Hybrid orbitals
Resonance
Low-level unfilled d-orbitals
AXmEn notation (X means ligands, m is the number of ligands, E is lone pair, n is the number of lone pairs, A is the central atom)
II. Concepts
The key to understanding bonding is realizing that bonds form because the products of the reaction have a greater stability than the atoms from which they form. For example, a solitary oxygen atom has only six electrons in its valence shell; however, if it forms a double bond with another oxygen atom (one sigma and one pi bond) then both will have a full valence shell of eight electrons and become isoelectronic to Neon, a noble gas. In general, bonds form so that the atoms involved can become isoelectronic to a noble gas. In every case, the atoms involved have less energy after the bond is formed than they had before. Therefore, bond formation is always exothermic. Bond destruction is always endothermic.
In drawing a Lewis dot structure, the octet rule is obeyed by most atoms. All Ligands obey the rule of octets, except hydrogen which obeys the rule of duets. The central atom may violate the rule of octets. Elements that are in period three or higher may violate the rule of octets by using low level unoccupied d orbitals for bonding. An example of this is sulfur hexafluoride, which has 12 electrons around the central atom and uses two 'd' orbitals along with one 's' orbital and three 'p' orbitals to form six identical hybrid orbitals with which it can bond. Lewis dot structures for polyatomic ions must place the extra electrons (or remove the lost electrons) from the central atom.
Electrons all have negative charges and therefore repel each other. The greater the concentration of negative the greater the strength of the repulsions. It is similar to kitchen magnets, the larger the magnet the more it repels other magnets. The greatest concentration of negative charge is in a lone pair and therefore it causes the greatest repulsions and therefore the greatest distortions of the shapes of the molecules.