Intro to ochem Flashcards
(170 cards)
1
Q
sigma bonds
A
forms when the bonding pair of electrons are localized to the space directly between the two bonding atoms. it is the lowest energy strongest and most stable type of covalent bond
2
Q
a pi orbital
A
is created by overlapping p orbitals
3
Q
sigma bonds are formed
A
in the area where hybrid orbitals of two atoms overlap
4
Q
pi bonds are formed
A
by the overlap of pure p orbitals
5
Q
number of hybrid orbitals must equal the
A
number of sigma bonds plus the lone pairs
6
Q
valence shell electron pair repulsion
A
states that electrons in an orbital seek to minimize their energy by moving as far away from other electron pairs as possible, minimizing the repulsive forces between them
7
Q
sp angle
A
180
8
Q
sp2 angle
A
120
9
Q
sp3
A
109.5
10
Q
sp3d
A
90,120
11
Q
sp3d2
A
90,90
12
Q
delocalized electrons
A
sometimes bonding electrons are spread out over three or more atoms
13
Q
aromaticity
A
is the increased stability of a cyclic molecule due to electron delocalization. must follow huckel’s rule 4n+2 will be aromatic
14
Q
nucleophiles are also called
A
lewis bases, with a partial negative charge seeking positively charged nuclei
15
Q
electrophilic functional groups
A
have a partial positive charge and seek electrons
16
Q
structural isomer
A
have the same molecular formula but different bond-to-bond connectivity
17
Q
conformational isomers
A
are not true isomers. they are different spatial orientations of the same molecule
18
Q
highest energy conformation and lowest
A
eclipsed and staggered
19
Q
stereoisomers
A
two unique molecules with the same molecular formula and the same bond-to-bond connectivity
20
Q
enantiomers
A
are non-superimposable mirror images of one another. they have the same molecular formula and connectivity but are not the same molecule because they differ in their configuration. they must have opposite absolute configuration at each and every chiral carbon. have the same chemical properties
21
Q
chirality
A
a carbon is chiral when it is bonded to four different substituents
22
Q
relative configuration
A
is not related to absolute configuration but two molecules have the same relative configuration about a chiral carbon if they differ by only one substituent and the other substituents are oriented identically about the carbon
23
Q
to find how a configuration rotates light you have to find it
A
experimentally
24
Q
if the compound rotates plane polarized light clockwise
A
it is designated with + for dextrorotary
25
if the compound rotates plane polarized light counterclockwise
- levoratory
26
diastereomers
have the same molecular formula and same bond-to-bond connectivity but are not mirror images of each other and are not the same compound
27
diastereomers have
the same absolute configuration at one or more of their chiral centers.
28
how do diastereomers differ from each other
they differ in their physical properties rotation of plane-polarized light, melting points, boiling points, solubilities and their chemical properties
29
meso compound
has multiple chiral centers but is optically inactive they have a plane of symmetry through their center which divides into two halves that are mirror images of each other. considered achiral
30
epimers
are diastereomers that differ in configuraton at only one chiral carbon
31
anomers
are cyclic diastereomers that are formed when a ring closure occurs at an epimeric carbon.
32
cis/trans isomers
geometric isomers that have substituents on the same side (cis) or on a different side (trans)
33
cis molecules have
higher boiling points due to their substituents and lower melting point. they also have a dipole moment
34
stronger bases are strong
nucleophiles
35
alcohols boiling point property
higher molecular weight is a higher boiling point and goes down with branching
36
strongly electron donating groups
oxygen, alcohol, amine
37
moderately donating
-OR
38
weakly donating
-R
39
strong electron withdrawing groups
O-N=O R-N-R+ Cl-C-Cl
40
moderatly withdrawying
carbonyls, aledhydes, ketones, esters, carboxyilic acids, tosylates, cyanide
41
acid trend based on how strong
methyl>primary > secondary> tertiary
42
hydrocarbons do what
donate electrons
43
carbonyl groups do what
electron withdrawing
44
tosylates and mesylates are used as
protection of alcohols. the conversion to a sulfonate prevents the alcohol from acting as an acid or undergoing other reactions
45
ethers are very good
solvents, even better than alcohols
46
two important characteristics of carbonyls
1. planar stereochemistry
| 2. polarity: partial negative charge on oxygen partial positive charge on carbon
47
how are peptide bonds formed
the amino acid attacks the carbonyl carbon of the carboxylic acid on another amino acid creating an amide
48
least reactive and strongest carboxylic acid derivative
amides
49
aldehydes and ketones undergo what sort of nucleophilic reaction
addition due to not having a good leaving group
50
ketones and aldehydes exist as what at room temperature
as keto-enol tautomers in which there is a shift from a carbonyl to an alkene with an alcohol.
51
aldehydes and ketones react with alcohols forming
hemiacetals and hemiketals
52
hemiacetals and hemiketals reacted with alcohol results in
acetal and ketals
53
sucrose
1'1 glycosidic linkage: glucose and fructose or 1'2
54
maltose
alpha 1,4' glycosidic linkage two glucose molecules
55
lactose
beta 1'4 galactosidic linkage: galactose and glucose
56
cellulose
beta 1'4 glycosidic linkage a chain of glucose molecules
57
amylose (starch)
alpha 1'4 glycosidic linkage a chain of glucose molecules
58
amylopectin
alpha 1'4 glycosidic linkage a branched chain of glucose molecules with alpha 1'6 glucosidic linkages forming the branches
59
glycogen
alpha 1'4 glycosidic linkage: a branched chain of glucose molecules with alpha 1,6 glucosidic linkages forming the branches
60
extensive properties
are proportional to the size of the system
61
intensive properties
are independent of the size of the system
62
volume and moles are examples of
extensive properties
63
pressure and volume are examples of
intensive properties
64
the greater the translational kinetic energy of gas molecules
the higher the temp
65
dividing one extensive property by another gives
an intensive property
66
absolute zero
-273.15 degrees celsius
67
an increase in 1 degree celsius is equivalent to
one degree of kelvin increase
68
to go from celsius to kelvin
add 273
69
difference between kelvin and celsius
celsius is a relative scale while kelvin actually measures the thermal energy
70
open systems
can exchange both energy and mass
71
closed systems
mass cannot be exchanged but energy can be
72
isolated systems
energy and mass cannot be exchanged
73
state functions
is the physical condition of a system as described by a specific set of thermodynamic properties
74
examples of state functions
internal energy (U), temperature (T), Pressure (P), Volume (V), Enthalpy (H), entropy (S), Gibbs energy (G)
75
pathway functions
depend on the pathway used to achieve that state
76
examples of pathway functions
work and heat
77
thermal energy is energy on what level
macroscopic level
78
internal energy is energy on what level
microscopic level
79
internal energy includes
vibrational energy, rotational energy, translational energy, electronic energy, intermolecular potential energy, and rest mass energy
80
vibrational energy
is created by the vibration of atoms within a molecule. vibrational energy makes an insignificant contribution to internal energy. gas has no vibrational energy
81
rotational energy
is created by the rotation of a molecule around its center of mass
82
translational energy
is created by movement of the center of mass of a molecule
83
electronic energy
is the potential electrical energy created by the attractions between electrons and their nuclei. in a chemical reaction changing electronic energy accounts for the greatest change in internal energy
84
intermolecular potential energy
is created by intermolecular forces between molecular dipoles
85
rest mass energy
is the energy described by einstein's E=mc^2. the sum of these energies for a very large group of molecules is the internal energy.
86
internal energy of an ideal gas has to depends on
temperature
87
only two ways to transfer energy between systems
heat and work
88
zeroth law of thermodynamics
states that two systems in thermal equilibrium with a third system are in thermal equilibrium with each other
89
conduction
is thermal energy transfer via molecular collisions. it requires direct physical contact
90
an objects ability to conduct heat is called its
thermal conductivity (k)
91
convection
is thermal energy transfer via fluid movements. differences in pressure or density drive warm fluid in the direction of cooler fluid.ex ocean and air currents
92
radiation
is thermal energy transfer via electromagnetic waves
93
blackbody radiators
have an emissivity of 1 and absorb 100% of radiation energy
94
work done by the system is considered
negative
95
first law of thermodynamics
energy is conserved
96
energy is transferred out of the system in
in expansion delta E is negative
97
enthalpy equation
H=U (internal energy) +PV
98
enthalpy equation under constant pressure conditions
change in internal energy + P delta V
99
formation of bonds always requires what? and the breaking of bonds always results in?
energy, releasing energy
100
positive enthalpy change
endothermic
101
negative enthalpy change
exothermic
102
anabolic reactions are usually
endothermic
103
catabolic reactions are usually
exothermic
104
entropy change of the universe equals
change in entropy of the system+ change in entropy of the surroundings= change in entropy of the universe
105
how does temp affect entropy
it increases it
106
third law of thermodynamics
assigns a zero entropy value to any pure element or compound in its solid form at absolute zero and in internal equilibrium
107
change in entropy equation
q(heat)/ T
108
reactions under non-standard states free energy equation
delta g= delta g knot+ RTlnq
109
reactions at equilibrium gibbs free energy equation
delta g= -RTlnk
110
entropy units
J/K
111
if volume decreases in a given reaction then shift to
one with less gas molecules
112
relationship between atmospheres, mmhg and torr
1atm=760mmhg=760torr
113
internal energy equation
change in internal energy= heat transfer+ the work for a given physical process (E= Q+W)
114
adiabatic process internal energy equation
E=w
115
isothermal process internal energy equation
E=0 0=Q+W
116
isovolumetric process internal energy equation
w=0 E=q
117
STP
1 atm 273K and 22.4 liters
118
real gases behave most ideally when
temperature is high and volume is large
119
ideal gas
gas molecules have no size. zero molecular volume
gas molecules do not exert forces on one another
gas molecules have completely elastic collisions
the average kinetic energy of gas molecules is directly proportional to the temp of the gas
120
partial pressure
is the total pressure of the gaseous mixture multiplied by the mole fraction of the particular gas
121
real gases deviate from ideal gases when
high pressures and low temp
122
how real gases differ from ideal gases in pressure and volumer
Vreal>Videal
| Preal
123
specific heat units
J/kgK or cal/gC(degree celsius)
124
amount of energy required to raise one gram of water by one degree Celsius
1 Cal=1000cal=4184 J
125
coffee cup calorimeter
keeps the pressure constant therefore it measures heats of reactions
126
bomb calorimeter
a bomb calorimeter measures internal energy change in a reaction by keeping constant volume.
127
there is a negative slope on the phase diagram for water because
waters solid phase is less dense than its liquid phase
128
critical temperature is
the temperature above which the substance cannot be liquified regardless of how much pressure is applied
129
equivalents in an acid base reaction
is defined as the mass of acid or base that can donate or accept one mole of protons. ex: 1 molar solution of H2SO4 is called a 2 normal solution because it can donate 2 protons from each H2SO4
130
normality
equivalents/ Liter of solution
131
in solutions if the overall reaction releases heat then the new intermolecular bonds are
stronger than the intermolecular attractions within the pure substances
132
negative heat of solution
indicates the formation of stronger intermolecular bonds
133
positive heat of solution
indicates the formation of weaker intermolecular bonds
134
dissolution of one substance into another
causes an increase in entropy except for a gas into a liquid or solid
135
vapor pressure of the liquid
the pressure created by the molecules in the open space at equilibrium
136
when vapor pressure of a liquid is equal to atmospheric pressure
the liquid boils
137
the melting point is when the vapor pressure of the solid is equal to
the vapor pressure of the liquid phase of that substance
138
nonvolatile solute
a solute with no vapor pressure
139
volatile solute
a solute with a vapor pressure
140
nonvolatile solutes do what to vapor pressure
decrease it
141
volatile solutes do what to vapor pressure
increase it
142
for nonvolatile solution
use the sum of the partial pressure of pure liquid and the mole fraction of the liquid
143
partial vapor pressure for volatile solutions
Pv=XaPa+XbPb
144
solubility product
solids are left out of the equation
145
spectator ions
are ions not included in the equilibrium expression for solubility
146
nearly all ionic compounds containing nitrate, ammonium, and alkali metals
are soluble
147
ionic compounds containing halogens are
soluble except for mercury, silver and lead (Hg, Pb, Ag)
148
sulfate compounds are
soluble except for mercury, lead, and the heavier alkaline earth metals (Hg, Pb, Ca, Sr, Ba)
149
compounds containing the heavier metals Ca, Sr, Ba
are soluble when paired with sulfides and hydroxides
150
carbonates, hydroxides, phosphates, and sulfides
CO3, PO4, S, OH are generally insoluble
151
pressure increases the solubility of
gases
152
Henry's law describes the solubility of gases
C (solubility of gas a ) = Ka (henry's constant) Pv (vapor partial pressure of gas a above the solution)
153
gas solubility decreases with
an increase in temperature
154
heavier larges gases tend to
experience greater van der waals interactions and tend to be more soluble
155
gases that chemically react with the solvent
have a greater solubility
156
redox titrations are used for
finding the molarity of a reducing agent
157
half equivalence point
near the middle of the gradual increase
158
equivalence point
where the voltage suddenly shoots up
159
when choosing an indicator what do you look for
an indicator that changes color as close as possible to the expected equivalence point
160
electrical potential
the more positive the potential the more likely the reaction is to proceed
161
to find the electrical potentials
separate the reaction into its two halves and add the half reaction potentials
162
galvanic cell
offers an alternative pathway for the flow of electrons through phases. the electric potential generates a current from one phase to another in a conversion of chemical energy to electrical energy
163
where in the galvanic cell does oxidation and reduction occur
oxidation at the anode and reduction at the cathode
164
current travels in the opposite direction as
electrons
165
the galvanic cell always has positive
cell potential
166
electrolytic cell properties
have a negative emf and the cathode is negative while the anode is positive. needs an outside energy source to drive the reaction
167
half equivalence point
where the concentration of acid equals the concentration of its conjugate base
168
pH is equal to pKa at
the half equivalence point
169
PI for acidic or basic amino acids
PI is the average of the first two pka values for a basic amino acid and PI is the average of the second and third pka for the acidic amino acid
170
buffers are made from
weak acids and its conjugate base
