Chapter 10 Flashcards
(24 cards)
When is the system more stable, as potential energy gets more negative or positive?
The more negative, the more stable the system becomes
Lewis Theory
Positions of electrons in the molecule are NOT considered; quickest bonding theory to employ
Valence-bond Theory
Atomic orbitals overlap where electrons are shared
Atomic orbitals are averaged to give geometries (VSEPR)
Molecular Orbital Theory
Explains bond order, stability, magnetism, resonance
Ionic Compounds
Metal to non-metal ionic interaction
Covelant Bonds
Sharing of electrons between two nonmetal atoms
Octet Rule
Electrons are transferred or shared to give each atom a noble gas configuration
Coordinate Covalent Bond
When a single atom contributes both of the electrons to the bond
Writing Lewis Structures
- Find total # of valence electrons (add for anions and subtract for cations)
- Choose central and terminal atoms
- Draw two electrons between each pair of connected atoms
- Add lone pairs to outer atoms (except H) to give complete octets
- Place all remaining electrons in lone pairs on central atom
Formal Charges
Used to keep track of how electrons are shared
FC = # valence electrons in central atom - lone electrons - # bond pairs
Resonance
When two or more Lewis Structures are equally feasible
Exceptions to Octet Rule
Odd-electron species (Radicals), Incomplete Octets (Boron, Hydrogen), expanded valence shells (period 3 and up)
Polarity of Bonds
non-polar when the shared electrons are equally attracted to both atoms
Polar - shared electrons are more strongly attracted to an atom
Electronegativity
Ability of an atom to attract electrons in a covalent bond
Higher values for electronegativity indicate greater attraction
Non-metals have higher EN than metals
EN rises going right and up the periodic table
Electronegativity and Polarity of Bonds
If difference in EN is small = non-polar
If difference is intermediate = polar
If difference is large = ionic bond
Dipole Moment
Product of a partial charge and distance
Draw arrow towards the atom with the higher EN
How to assign VSEPR
- Draw Lewis Structure
- Count # of electron groups (each lone pair counts as 1, each bond group is 1 even if double)
- Predict molecular shape and state VSEPR
Electron and Molecular Geometry - Linear
Two charge clouds/electron groups Electron geom = Linear Molecular geom = Linear Bond angle = 180 VSEPR = AX2
Electron and Molecular Geometry - Trigonal Planar
Three charge clouds/electron groups *Three outer atoms + no lone pairs Molecular geom = Trigonal Planar Bond angle = 120 VSEPR = AX3 *Two outer atoms + one lone pair Molecular geom = Bent Bond angle = 120 VSEPR = AX2E
Electron and Molecular Geometry - Tetrahedral
Four charge clouds/electron groups *Four outer atoms + no lone pairs Molecular geom = Tetrahedral Bond angle = 109.5 VSEPR = AX4 *Three outer atoms + one lone pair Molecular geom = Trigonal-pyramidal Bond angle = 109.5 VSEPR = AX3E *Two outer atoms + two lone pairs Molecular geom = Bent Bond angle = 109.5 VSEPR = AX2E2
Electron and Molecular Geometry - Trigonal-bipyramidal
Five charge clouds/electron groups *Five outer atoms + no lone pairs Molecular geom = trigonal bipyramidal Bond angle = 90, 120 VSEPR = AX5 *Four outer atoms + one lone pair Molecular geom = see-saw Bond angle = 90, 120 VSEPR = AX4E *Three outer atoms + two lone pairs Molecular geom = T-shaped Bond angle = 90 VSEPR = AX3E2 *Two outer atoms + three lone pairs Molecular geom = Linear Bond angle = 180 VSEPR = AX2E3
Electron and Molecular Geometry - Octahedral
Six charge clouds/electron groups *Six outer atoms + no lone pairs Molecular geom = Octahedral Bond angle = 90 VSEPR = AX6 *Five outer atoms + one lone pair Molecular geom = square-pyramidal Bond angle = 90 VSEPR = AX5E *Four outer atoms + two lone pairs Molecular geom = Square-planar Bond angle = 90 VSEPR = AX4E2
Bond Order
Single bond = 1
Double bond = 2
Triple bond = 3
*Bond length decreases with increasing bond order
Bond Dissociation Energies
Energy required to break one mole of covalent bonds for gaseous species
- Energy is release when bonds form; energy absorbed when bonds break
- Use changes of enthalpy reactions to calculate energy for bonds (Hess’ Law)
