Chapter 3 - Bonding and Chemical Interactions

incomplete octet

elements are more stable with

expanded octet

elements are more stable with >8 e- in their valence shell
phosphorus (10)
sulfur (12)
chlorine (14)
etc

ionic bonding

one + electrons from an atom with low ionization energy (most likely a metal) transferred to an atom with high electron affinity (typically a non metal), the resulting electrostatic attraction between opposite charges is what holds the ions together, and no electrons are shared between atoms (SIGNIFICANT DIFFERENT ELECTRONEGATIVITY)

colavent bonding

electron pair is shared between two atoms, typically have similar values of electronegativity
\*\*coordinate covalent if the shared electrons are contributed by only one of the atoms

further in electronegativity the more ionic character

covalent compound properties

low melting/boiling points

bond energy

the greater the number of pairs of electrons shared between the atomic nuclei, the more energy required to break the bonds holding the atoms together

polarity

occurs when two atoms have relatively different electronegativities, the atoms with higher electronegativity gets the larger share of electron density
- higher electronegativity, partial NEGATIVE
- lower electronegativity, partial POSITIVE

non polar covalent bond

there is no separation of charge across the bond, difference in electronegative >0.5

dipole moment

vector quantity of a polar covalent moment, electrons pulled more towards the atoms with higher electronegativity (giving it partial negative)

formal charge equation

valance e - nonbonding (dots) - 0.5*bonding (lines)

VSEPR model, 2 regions of e- density

linear, 180 degrees

VSEPR model, 3 regions of e- density

trigonal planer, 120 degrees

VSEPR model, 4 regions of e- density

tetrahedral, 109.5 degrees

VSEPR model, 5 regions of e- density

trigonal bipyramidal, 90/120/180 degrees

VSEPR model, 6 regions of e- density

octahedral, 90/180 degrees

molecular orbit

two atoms bond to form a compound and the atomic orbitals interact

sigma bonds

orbitals overlap head to head, linear, free rotation

pi bond

orbitals overlap to create two parallel electron cloud densities, no free rotation

van der Waals force

attractive or repulsive interactions of short lived and rapidly shifting dipoles created from rapid polarization and counter polarization of electron cloud density

* weakest of all intermolecular forces
* stronger in larger molecule and molecules of the same shape
* hydrocarbons typically participate in h

diople-dipole interactions

present in liquid and solid phases
polar molecules tend to orient themselves that the partial charges of each dipole align opposite (ie partial postive next to partial negative)
- tend to lead to higher boiling/melting points than non polar bonds

hydrogen bonds

unusually strong dipole dipole interaction
no sharing/transfer of electrons
H bonding to N, O, F has virtually no electron density, making it a prime target to interact with partial negative of N, O, F
unusually high boiling point

polarity of dipole moment equation

p=qd

q = charge

d = distance

coordinate covalent bond

When both electrons shared in bond are donated by one atom

transition metals can make these type of bonds

also known as divalent bond