Ch. 3: Bonding Flashcards
what makes carbon so special?
- carbon has unique bonding properties
- it is tetravalent (it can form bonds with up to four other atoms, allowing for the massive versatility required to form the foundation of biomolecules and life itself)
- as it is located near the center of the periodic table, it can form bonds with many different elements because of its moderate electronegativity
- because carbon atoms are fairly small, the bonds that they form are strong and stable
defn: ionic vs. covalent bonds
ionic = electrons are transferred from one atom to another and the resulting ions are held together by electrostatic interactions
covalent = electrons are shared between atoms
defn + range: principal quantum number (n)
corresponds to the energy level of a given electron in an atom and is essentially a measure of size
the smaller the number, the closer the shell is to the nucleus and the lower its energy
MCAT tests on 1-7
defn + range: azimuthal quantum number, l
describes the subshells that can be within each electron shell
range: 0 - n-1 for a given energy shell
what do the l values 0, 1, 2, and 3 correspond to?
the s, p, d, and f subshells
defn + range: magnetic quantum number, ml
describes the orbitals within each subshell
range: - l to l
char: s-orbital
spherical and symmetrical, centered around the nucleus
char: p-orbital
composed of 2 lobes located symmetrically about the nucleus and contains a node (an area where the probability of finding an electro is zero) at the nucleus
you can picture it as a dumbbell that can be positioned along the x, y, or z axis
char: d-orbital
composed of 4 symmetrical lobes and contains 2 nodes
four of the d-orbitals are clover-shaped and the fifth looks like a donut wrapped around the center of a p-orbital
defn + range: spin quantum number, ms
each orbital can hold two electrons, which are distinguished by this quantum number
the only values are -0.5 and 0.5
defn: molecular orbitals
when two atomic orbitals combine, this is what they form
how are molecular orbitals obtained mathematically?
by adding or subtracting the wave functions of the atomic orbitals
defn: bonding orbital vs antibonding orbital
BONDING: produced if the signs of the wave functions are the same; lower-energy, more stable
ANTIBONDING: produced if the signs of the wave functions are different; higher-energy, less stable
defn: sigma bond
results when a molecular orbital is formed by head-to-head or tail-to-tail overlap
defn: pi bond
forms when two p-orbitals line up in a parallel (side-by-side) fashion and their electron clouds overlap
this is a bonding molecular orbital
are all single bonds sigma bonds or pi bonds?
sigma bonds
how does a double bond form from sigma and pi bonds? what about a triple bond?
double bond: one pi bond on top of an existing pi bond
triple bond: a sigma bond and 2 pi bonds
can a pi bond exist independently of a sigma bond?
no, only after the formation of a sigma bond will the p-orbitals of adjacent carbons be parallel and in position to form the pi bond
why are shorter bonds stronger than longer bonds?
they hold atoms more closely together and require more energy to break
double bonds are stronger than single bonds, but are individual pi bonds stronger than sigma bonds?
no, individual pi bonds are weaker than sigma bonds
thus, it is possible to break only one of the bonds in a double bond, leaving a single bond intact
what is the difference in freedom of motion between single and multiple bonds?
single bonds: allow free rotation of atoms around the bond axis
double and triple bonds: hinder rotation, and in effect, lock the atoms in position
the more bonds that are formed, the shorter or longer the bond length?
the more bonds = the shorter the bond length
func: hybridization
a way of making all of the bonds to a central atom equivalent to each other
how are hybrid orbitals formed?
formed by mixing different types of orbitals
char (3): sp3 hybridization
- use advanced math to merge 3 p-orbitals and 1 s-orbital, which forms 4 identical sp3 orbitals with new, hybridized shapes
- all 4 of these orbitals point toward the vertices of a tetrahedron to minimize repulsion, which explains the tetrahedral geometry
- there are no unhybridized p-orbitals to form pi bonds
how is sp3 hybridization accomplished?
by promoting one of the 2s electrons into the 2pz orbital
this produces 4 valence orbitals, each with one electron
how would you answer an MCAT question that tests how much “s character” a certain hybrid orbital has, for example?
we simply need to determine the type of hybridization and use the name to solve
in sp3, we have 1s and 23p, so the bond has 25% s and 75% p character
what is the hallmark shape and hybridization of carbon-containing compounds?
is this the only possibility for carbon?
sp3, tetrahedral
no!
char (5): sp2 hybridization
- 1 s-orbital is mixed with 2 p-orbitals to make 3 sp2-hybridized orbitals
- seen in alkenes
- the third p-orbital of each carbon is left unhybridized, this is the orbital that participates in the pi bond
- the 3 sp2 orbitals are oriented 120deg apart, allowing for maximum separation
- trigonal planar geometry
what do the hybrid orbitals in sp2 hybridization do?
for example, in ethene
2 of the sp2 hybridized orbitals will participate in C-H bonds and the other hybrid orbital will line up with the pi bond to form the sigma component of the C=C double bond
char (4): sp hybridization
- seen in triple bonds
- 2 of the p-orbitals form pi bonds, the third p-orbital will combine with the s-orbital to form two sp-orbitals
- these orbitals are oriented 180deg apart, which explains the linear structure of molecules containing sp-hybridized carbons
- the two pi bonds can be between the carbon and one other atom (forming a triple bond) or between the carbon and two different atoms (forming 2 double bonds in a row)
when does resonance delocalization of electrons occur?
in molecules that have conjugated bonds
process (2): conjugation
- requires alternating single and multiple bonds because this pattern aligns a number of unhybridized p-orbitals down the backbone of the molecule
- pi electrons can then delocalize through the p-orbital system, adding to stability of the molecule
explain why resonance structures aren’t in equilibrium?
the electron density is distributed throughout and the true form is a hybrid of the resonance structures
unless the stability of the various resonance forms differ, then the true electron density will favor the most stable form
what are four reasons why particular resonance structures might be favored?
- they lack formal charges
- they form full octets on highly electronegative atoms
- there is stabilization of positive and negative charges through induction
- aromaticity
