1. BONDING Chapter 6

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

3. I ONIC B ONDING metal (cation) + nonmetal (anion) electrons are gained or lost

4. C OVALENT B ONDING nonmetal + nonmetal electrons are SHARED nonpolar – electrons are shared equally (examples: F 2 and O 3 polar – unequal sharing of electrons (examples: H 2 O and NH 3 )

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. C HARACTERISTICS OF C OVALENT B ONDS 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. L EWIS D OT S TRUCTURES Octet rule – compounds are formed so that each atom has eight electrons in its outer shell EXCEPTIONS – hydrogen & helium Examples: N, H, F, F 2 and NH 3

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: O 2 and C 2 H 4 ) triple bond – sharing three pairs of electrons between two atoms (examples N 2 and C 2 H 2 )

9. R ESONANCE More than one correct Lewis dot structure SO 2 SO 3

10. I ONIC B ONDING 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. P OLYATOMIC IONS nitrate = NO 3 - nitrite = NO 2 - sulfate = SO 4 -2 sulfite = SO 3 -2 hydroxide = OH - ammonium = NH 4 + carbonate = CO 3 -2 bicarbonate = HCO 3 -

12. M ETALLIC B ONDING chemical attraction that results from the attraction of metal atoms and surrounding electrons high electrical and thermal conductivity malleability ductility

13. M OLECULAR G EOMETRY 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 (F 2 ) or central atom bonded to two other atoms with NO lone pairs (CO 2 ) Bent – central atom bonded to two other atoms and has lone pairs (water)

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

15. B OND A NGLES linear = 180 o bent = 120 o (when there is 1 lone pair) bent = 109 o (when there are 2 lone pairs) trigonal planar = 120 o trigonal pyramidal = 109 o tetrahedral = 109 o

16. I NTERMOLECULAR F ORCES 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 (O 2 ) – also called van der Waals forces