 |
| Chapter 8 |
| Electronegativity- how stongly an atom attracts electrons F has the highest electronegativity A metal with a non-metal automatically equals an ionic bond An element minus itself is covalent <.4 = covalent .5-1.9 = polar bonds >2.0 = Ionic bonds E= 2.31x10^-19 J.nm (Q1 Q2 / r) E= energy r= distance between ion centers (nm) Q= charge Dipoles= polar Ionization energy increases as is goes right across the periodic table and decreases down. Electronegativity increases as it goes right and decreases as it goes down. The size of the atoms decreases as it goes right and increases as it goes down. |
|
| XeF2, BeCl2, CO2, N2O are all Linear molecules Bent molecules are always polar. NH3, PCl3 are both pyrimidal. |
|
| VSEPR - theroy ( valence shell electron pair repulsion theroy) *Lone pair electrons exert a stronger repulsive effect on adjacent electron pairs then do bonded pairs. *Multiple bonds do not effect the gross strong chemistry of the molecule. *Electron pairs in the valence shell of an atom are arranged around that atom in such a way as to minimize the repulsion. *The core electrons do not influence the shape of the molecule. |
|
| Delta H = Sumation of D bonds broken minus the sumation of D bonds formed H=Ed(bb)-Ed(bf) ex) Calculate Delta H for: CH4 + 4F2 --> CF4 + 4HF Bonds Broken: Bonds Formed: 4 C-H 413 x 4= 1652 4 C-F 485 x 4= 1940 4 F-F 154 x 4= + 616 4 H-F 515 x 4= 2260 2268 KJ 4200 KJ Delta H = 2268 - 4200 = -1932 KJ (exothermic) |