Use LEFT and RIGHT arrow keys to navigate between flashcards;
Use UP and DOWN arrow keys to flip the card;
H to show hint;
A reads text to speech;
27 Cards in this Set
- Front
- Back
|
Exception to octet rule: incomplete octet |
Elements that are stable with fewer than eight electrons in their valence shell. Hydrogen (2), lithium (2), Helium (2), beryllium (4), Boron (6) |
|
|
Exception to octet rule: expanded octet |
any element in period 3 and greater can hold more than eight electrons, including phosphorus (10), sulfur (12), and chlorine (14) |
|
|
Exception to octet rule: odd number of electrons |
any molecule with an odd number of valence electrons cannot distribute those electrons to give eight to each atom; for example, nitric oxide (NO) has eleven valence electrons |
|
|
Common elements that abide by octet rule |
C, N, O, F, S, Mg |
|
|
Ionic bond |
one or more electrons from an atom with a low ionization energy, typically a metal, are transferred to an atom with a high electron affinity, typically a nonmetal (occurs when difference in electrongegativity is 1.7 or greater). Resulting electrostatic attraction holds the ions together. High melting point, good conductors when dissolved. Crystalline lattice. |
|
|
Covalent bond |
An electron pair is shared between two atoms, typically nonmetals, that have relatively similar values of electronegativity. Can be polar (if difference is between 0.5 and 1.7 or greater than 1.7 and two nonmetals) or nonpolar bond (if difference is less than 0.5) |
|
|
Coordinate covalent |
If both of the shared electrons are contributed by only one of the two atoms. Ex: Lewis base and Lewis acid pair |
|
|
Bond order |
the number of shared electron pairs between two atoms |
|
|
Lewis structure |
Most stable structure reduces number and magnitude of formal charges. Less separation between opposite charges is preferred. Negative formal charges placed on more electronegative atoms is more stable. |
|
|
Drawing Lewis structures |
Draw out backbone (least electroneg. atom is usually central atom). Count all the valence electrons, this is the number of valence electrons of the molecule. Draw single bonds. Complete the octets. Place any extra electrons on the central atom. |
|
|
Formal Charge |
formal charge = V - Nnonbonding - 0.5 Nbonding |
|
|
Difference between formal charge and oxidation number |
formal charge minimizes effects of electronegativity while oxidation number maximizes. Reality is between the two. |
|
|
VSERP Theory |
Draw the lewis structure of the molecule. count the total number of bonding and nonbonding pairs in the valence shell of the central atom. arrange the electron pairs around the central atom so that they are as far apart as possible. Make sure you make distinction between molecular and electronic geometry |
|
|
Linear |
180 |
|
|
Trigonal planar |
120 |
|
|
Tetrahedral |
109.5 |
|
|
Trigonal bipyramidal |
90, 120, 180 |
|
|
octahedral |
90, 180 |
|
|
Electric geometry |
describes the spatial arrangement of all pairs of electrons around the central atom, including both the bonding and nonbonding electrons |
|
|
Molecular geometry |
describes the spatial arrangement of only the bonding of electrons |
|
|
Correction to Ideal bond angle |
Tetrahedral electronic geometry is associated with an ideal bond angle of 109.5, but nonbondingpairs able to exert more repulsion than bonding pairs because these electrons reside closer to the nucleus. Thus, the angle in ammonia is closer to 107 and the angle is water is closer to 104.5. |
|
|
bonding orbital |
if the signs of the two overlapping atomic orbitals are the same |
|
|
antibonding orbital |
if the signs of the two overlapping atomic orbitals are opposite |
|
|
Sigma bond |
when orbitals overlap head to head. Allows free rotation about the bond |
|
|
Pi bond |
when orbitals overlap in such a way that there are two parallel electron cloud densities. Due to this parallel nature, there is no free rotation about the bond |
|
|
London dispersion forces
|
attractive interactions of short-lived and idly shifting dipole moments. a type of van der waals force |
|
|
Dipole-dipole interactions |
become negligent in the gas phase because of distance between molecules |
