Valence Bond Theory

How do chemical bonds work

When two bonds get closer together, the nucleus attracts electrons from the other nucleus (decreasing potential energy) but if they get too close, the nuclei repel each other.

Valence Bond theory

A chemical bond results from the overlap of 2 orbitals, each half filled

If they do not overlap, they repel

pi bond

a covalent bond with electron density on either side of internuclear axis

sigma bond

covalent bond with electron density along the axis

bond hybridization

mixing different types of atomic orbitals to produce a set of equivalent hybrid orbitals for boding.

Unpaired electrons/ left over p orbitals

occupies pi molecular orbitals. if there is 2 unoccupied p orbitals, they will be perpendicular to each other

Bonding Axis in linear molecules

z-axis

Bonding axis in plane molecules

z-axis is perpindicular (x or y)

What happens when a base reacts with an acid?

The energy of the highest occupied orbital gets lowered

Adjacent p orbitals that are parallel

pi bond

triple bond

(2pz is sigma bond) 1 sigma, 2 pi

What happens when a hybrid orbital has more p character

Energy is raised, decreased bond angle, forms electron withdrawing atom

What happens when a hybrid orbital has more s character

Lowers energy, is good for lone pairs

sp

180

sp2

120

sp3

109

p

90

Functional Group

a group of atoms making up a portion of a molecule that have a characteristic structure or reactivity

Alklanes

simplest group, composed only of C and H, only sigma bonds, carbon is tetrahedral

Van Der Waals interactions

a term to describe intermolecular forces that result from dipoles between molecules

London forces

intermolecular forces involving instantaneous and induced dipoles.

Magnitude of attraction and length of dipole

Larger molecule → larger dipole → larger attraction

Not chiral

if molecule has a mirror plane

Why is p character better for electron withdrawing atoms?

p orbitals are higher in energy and in bonds, more likely to lose electron from its higher energy orbital so it’s more likely for an electron to go to another high energy orbital as well

Lewis Acid

electron pair acceptor (low energy empty orbital to accept electrons)

Lewis Base

electron pair donor (high energy electrons you can donate that aren’t being shraed)

Bronstead acid

a proton donor (Accepts protons and electrons)

What does lower energy orbitals have to do with H+

When H+ is removed, there is lower energied electrons which makes it more stablized

Staggered

atom on one side of bond try to get in between atoms in the back (minimizes repulsion between electron densities)

Eclipse

front hydrogens are covering the back ones ( you can’t see them anymore)

Photon added moving electron in pi

would be moved to antibonding pi and would now be able to rotate

Formal Charge gives insight to

dipole moment

Site of unsaturation

(a ring or a pi bond) a place on a molecule where you can add additional atoms (without violating octet rule)

SOU Equation

(2C + N - H - X + 2)/2

SOU Examples

Alkyne would have 2 SOU because it has two pi bonds, an alkene has 1 SOU

For every SOU…

2 hydrogens are lost

When there is no SOU…

for every 1 carbon, 2 hydrogens are added

Cyclic molecules

highly strained because of small bond angles and eclipsed hydrogens (all angles are 109 degrees)

Isomers

compounds with the same molecular formula but with different atomic structures