Exam 1 Flashcards
What is the limit on the number of valence electrons the second row of elements can have?
8
formal charge main formula and condensed
Main: Formal charge=(#of starting valence electrons)-((lone pair e)+(1/2 bonding e))
or
valence electrons-(Dots+lines)
covalent bond
Electrons are shared equally between two atoms. Aka nonpolar bond, electronegativity difference is between 0 and 0.4
Polar covalent bond
electrons are shared between atoms but not equally, resulting in induction. Electronegativity difference is between 0.4 and 2
induction
withdrawal of electrons toward an atom, resulting in partial negative and positive charges
Ionic bond
Electrons are not shared, but form a bond due to the force of attraction between opposite charges. Electronegativity difference is between 2 and 4. No electron charge, purely electrostatic.
Which atom is an exception to the octet rule and why?
H because it doesn’t have enough electron orbitals
What happens when atoms obey the octet rule?
atoms are stabliized at a lower energy level
What formal charge do neutral atoms have?
0
what does a formal charge tell you?
how far the charge is from the correct number of valence electrons
what does electronegativity indicate?
how strongly atoms pull electrons towards themselves
Vectors
Arrows in lewis diagrams pointing the more electronegative and more electron dense atoms. Length of vector indicates magnitude
dipole moment
quantification of the induction that occurs in polar covalent bonds
molecular bond polarity
the vector sum(or middle of the vectors) of a molecule with more than 2 atoms. This describes what the entire molecule is doing
intermolecular forces vs intramolecular forces
inter: BETWEEN molecules, not true bonds and weaker than true bonds
intra: with the same molecule
intermolecular forces from strongest to weakest
H bonding, dipole-dipole then dispersion forces
how to tell strength of dispersion forces
more surface area(#of C atoms) means stronger dispersion forces because charges are farther away from each other
How does branching effect surface area?
branching decreases it
atomic orbitals
describe where electrons exist within an atom based on wave equations that take into account the wave like nature of electrons, refers to space with 90-95% electron density
what happens with orbitals in bonds between atoms?
orbitals must overlap so electrons can be shared
S orbital
spherical in shape, almost all atoms have them
P orbitals
“dumbell” or infinity sign shaped
Aufbau principle
fill lowest energy orbital first
Pauli
each orbital accomodates max of 2 electrons with opposite spin
Hund’s rule
one electron is place in each orbital before any electrons are paired
degenerate orbitals
orbitals with the same energy level
Molecular orbital theory
(MO theory) a sophisticated approach to visualize atomic orbital overlap
Molecular orbitals
when electrons are shared between two atoms, their 2 atomic orbitals become molecular orbitals. Number of AO always = number of MO
antibonding MO
formed from destructive interference of waves, higher than bonding orbital in energy, make nodes
Bonding MO
formed from constructive interference of waves, lower than antibonding orbital in energy, make amplification, fills up before bonding MO because lowest energy orbitals are filled up first
why do nodes give antibonding MO’s higher energy
they prevent sharing of electrons, which stabilizes atoms
node
area where there is 0% of electrons existing
Highest energy occupied molecular orbital
(HOMO) highest energy orbital among the occupied orbitals, bonding orbital
lowest energy unoccupied molecular orbital
(LUMO) lowest energy orbital among the unoccupied orbitals, antibonding MO
Why is it important to know HOMO and LUMO
they react and share electrons
What to solid wedge-shaped bonds and dashed bonds indicate in illustrations?
solid-wedge shaped bonds are coming towards you and dashed bonds are going into the drawing away from you
Hybridization
a solution to the problem that MO theory is too complex to analyze without computers. It involves mentally averaging the effects of valence orbitals, giving hybridized orbitals that are equal in energy and unhybridized orbitals
P bond physical properties compared to S
P bonds are longer and weaker and have more energy
sp3 orbitals
4 hybridized orbitals that are equivalent in energy, tetrahedral geometry, 25% s character and 75% p character
sp2 orbitals
33% s character, 66% p character, trigonal planar geometry, 3 hybridized sp2 orbitals(equal in energy) and 1 unhybridzed p orbital.
sigma bonds
bonds formed by head to head overlap of hybridized orbitals, experience free rotation
pi bonds
formed by side-by-side orbital overlap of unhybridized p orbitals, dont experience free rotation so they maximum orbital overlap
sp orbitals
have 2 hybridized and 2 unhybridized orbitals. 50% s and 50% p. Linear geometry. Shortest and strongest bonds of all hybridized orbitals.
Why do s orbitals have shorter bonds?
S orbitals have spherical shapes
steric number
indicates number of electron pairs(both bonding and non-bonding-lone pairs) that are repelling each other. = # of sigma bonds + number of lone pairs
what is the trend with bond strength and stability?
direct
what do line represent in bond line structures?
covalent bonds
heteroatoms
atoms that aren’t C or H in bond line structures
how are triple bonds drawn in bond line figures?
straight lines with surrounding single bonds
what is the first step to drawing the bond line structure of a complex molecule?
look for the longest continuous carbon chain excluding rings of carbon
substituents
groups coming off the longest carbon chain in a bond line structure
functional groups
chemical arrangements of certain atoms that undergo specific, predictable reactions
simplest functional group
alkanes
IR spectroscopy
stands for infra red spectroscopy, utilizes photons of light of a certain energy that cause transitions in the vibrational energy levels of atoms. Functional groups and bonds have their own, unique signals as energy peaks in the spectroscopy graph
which common heteroatom is an exception to the octet rule?
sulfur
HONC acronym
HONC
1 2 3 4
#of bonds matched to element
carbocation
+1 positive carbon, missing H/proton
carboanion
-1 charged carbon, three bonds with a lone pair
what does formal charge tell you about carbons in the bond line structure?
it allows you to figure out how many lone pairs and atom has, formal charge combined with normal # of valence electrons tells you how many dots and lines total go around an atom
which orbitals can overlap with each other?
unhybridized p-orbitals
what is the inadequacy of bond line structures?
they aren’t good for showing reasonance and make it look like electrons are stuck in between 2 carbons with unhybridized p-orbitals when they actually aren’t
resonance
sharing of electron density throughout and overlapping pi p-orbital system
original structure
first structure showing a resonance structure
contributor
other structures that are apart of the resonance structures, apart from the contributor
resonance method
a method chemists used to deal with the inadequacy of bond line structures, it represents the distribution of electrons throughout a molecule’s pi orbitals
delocalization
the spreading of a charge, makes a molecule more stable and have less energy because electrons are spread throughout the whole molecule, reducing reactivity, aka resonance stabilization
which bond do and dont affect resonance?
pi bonds only affect resonance, sigma bonds do not
what does a double bond represent with the orbitals and atoms it is connecting?
it represents one p orbital at each atom that is sharing electrons with the other p orbital
5 patterns for resonance structure resonance
- allylic lone pairs
- allylic positive charge
- lone pair of electrons adjacent to a positive charge
- a pi bond between two atoms with different electronegativities
- conjugated pi bonds in a ring
vinyl
atoms directly bonded to a pi bond
allyl
atoms one atom away from pi bond
how to address an allylic lone pairs
2 arrows: one from the lone pair to the bond next to the pi bond, and another from the pi bond to the lone pair on the other side of the allyl
how to address an allylic positive charge
draw and arrow from the pi bond to the bond with the positive charge
how to address a lone pair adjacent to a positive charge
draw an arrow from the lone pair to the bond with the positive charge
how to address a pi bond between 2 atoms of differing electronegativity
draw an arrow from the pi bond to the more electronegative atoms
Conjugated pi bond in a ring pattern
atom atom must have unhybridzed P orbitals that overlap with its neighbors(alternating pi and single bonds). you can move them clockwise or counterclockwise in the ring
Rules for determining the most important resonance structure
Rules by priority:
- the most significant structure has the greatest number of filled octets
- the structure with fewer formal charges is more significant
- A structure with a negative formal charge on the more electronegative element will be more significant
- structures with a carbo cation are usually less stable
localized electrons
not in resonance, in hybridized orbitals, dont move, decreases stability
delocalized electrons
in resonance, in unhybridized p orbitals, increase stability, move around the whole structure, occur in sp or sp2 bonds
Bronsted-Lowry definition of acids and bases
Acid: donate a proton(deprotonation)
Base: Bases accept a proton(protonated)
Conjugate acids and bases
Conjugate acid: results when a base accepts a proton
Conjugate base: results when an acid gives up a proton
how many steps are all acid/base reactions?
1
physical characteristics of strong acids
a strong acid is unstable and will react quickly to form a stable conjugate base
Ka
equilibrium constant that gives us an idea of how many reactants are turning into products, values range from 10^-50 to 10^10
pKs
-log(Ka), range from -10 to 50, exponentially related to Ka
what do different numerical values of Ka represent
Ka>1: more reactants are going to products and vice versa
