1. BONDING
Unit 6

2. Chemical Bonding
Chemical bond – mutual electrical attraction between the nuclei and valence electrons of different atoms that bind atoms together

3. Ionic Bonding metal (cation) + nonmetal (anion)
electrons are gained or lost

4. Covalent Bonding nonmetal + nonmetal electrons are SHARED
nonpolar – electrons are shared equally (examples: F2 and O3
polar – unequal sharing of electrons (examples: H2O and NH3)

5. molecule – covalent compound (a neutral group of atoms that are held together by covalent bonds)
chemical formula – number of atoms in a compound (example: water)

6. Characteristics of Covalent Bonds
bond length – the distance between two bonded atoms
small bond length = strong bond
long bond length = weak bond
bond energy = energy required to break a bond
high energy = strong bond
low energy = weak bond

7. Lewis Dot Structures
Octet rule – compounds are formed so that each atom has eight electrons in its outer shell
EXCEPTIONS – hydrogen & helium
Examples: N, H, F, F2 and NH3

8. unshared electrons are called lone pairs
single bond – sharing one pair of electrons between two atoms
double bond – sharing two pairs of electrons between two atoms (examples: O2 and C2H4)
triple bond – sharing three pairs of electrons between two atoms (examples N2 and C2H2)

9. Resonance
More than one correct Lewis dot structure
SO2
SO3

10. Ionic Bonding cations and anions form a neutral compound
lattice energy – the energy released when one mole of ionic crystalline is formed from gaseous ions
ionic bonds are stronger than covalent
covalent compounds tend to have lower MP and BP
ionic compounds are hard and brittle

11. Polyatomic ions nitrate = NO3- nitrite = NO2- sulfate = SO4-2
sulfite = SO3-2
hydroxide = OH-
ammonium = NH4+
carbonate = CO3-2
bicarbonate = HCO3-

12. Metallic Bonding
chemical attraction that results from the attraction of metal atoms and surrounding electrons
high electrical and thermal conductivity
malleability
ductility

13. Molecular Geometry
VSEPR theory – repulsions between the sets of valence electrons around the atoms causes the sets to be as far apart as possible
SHAPES:
Linear – only two atoms bonded (F2) OR central atom bonded to two other atoms with NO lone pairs (CO2)
Bent – central atom bonded to two other atoms and has lone pair(s) (water & SO2)

14. Trigonal planar – central atom bonded to 3 other atoms with NO lone pairs (SO3)
Trigonal pyramidal – central atom bonded to 3 other atoms and 1 lone pair (NH3)
Tetrahedral – central atom bonded to 4 other atoms with NO lone pairs (CH4)

15. Intermolecular Forces
dipole-dipole – between two different elements in the same compound (H- Cl)
hydrogen bonding – strong dipole forces between H and SONF (the strongest IMF) (water and ammonia)
London dispersion forces – between the same type of atom (O2) – also called van der Waals forces