Lecture 8 Flashcards
1
Q
metals
A
fewer valence electrons, resulting in lower IE, therefore a lower electronegativity
2
Q
non-metals
A
more valence electrons, resulting in higher IE, therefore a higher electronegativity
3
Q
ionic bonding
A
full transfer of electrons from a metal to a non-metal
4
Q
covalent bonding
A
equal sharing of 2 valence electrons in a bonding orbital between atoms of similar electronegativities, therefore no polarization
5
Q
polar covalent bonding
A
- uneven sharing of 2 valence electrons in a bonding orbitals due to difference in electronegativies
- stronger because of the opposite charges
6
Q
metallic bonding
A
- valence electrons delocalized over many atoms via their valence orbitals (bonds of orbital)
- delocalization of electrons leads to conductivity, thermal conductivity, or ductility
7
Q
covalent bond strength
A
- depends on overlap between the bonding atomic valence orbitals
- orbitals must point towards each other
- same size orbitals cause better overlap
8
Q
ionic bond strength
A
- small, medium sized orbitals cause good overlaps
- one large and one small orbital causes poor overlap
- both large orbitals causes decent overlap
9
Q
Valence Shell Electron Pair Repulsion (VSEPR)
A
bonding pairs and lone pairs repel and molecules will adopt geometries that minimize repulsions by maximizing space
10
Q
AX2
A
- linear
- 180°
11
Q
AX3
A
- trigonal planar
- 120 °
12
Q
AX4
A
- tetrahedral
- 109.5°
13
Q
AX5
A
- square pyramidal
- trigonal bipyrimidal
14
Q
AX6
A
- octahedral shape
- 90°
