CHEMISTRY/ORGANIC CHEMISTRY FINAL REVIEW Flashcards
Atomic number
Z - number of protons found in an atom of that element
Mass number
A - sum of the protons and neutrons in the atom’s nucleus
-varies in isotopes
Planck relation
E = hf
Energy is related to frequency times Planck’s constant
Principal quantum number, maximum number of electrons per number
n - indicates the electron’s shell
Maximum number of electrons within a shell = 2n^2
Azimuthal quantum number, range of possible values
l - shape and number of subshells
range of possible values: 0 to n-1
- only one subshell in first principal energy level (0)
- two in second principal energy level
- three in third principal quantum level
indicated as a letter (s,p,d,f)
Maximum number of electrons within a subshell = 4l + 2 (s2,p6,d10,f14)
Magnetic quantum number, range of possible values
ml - specifies the electrons orbital
range of possible values: -l to l
orbitals in s are spheral, p are dumbbell
Spin quantum number
ms - designated +½ or -½
Electron configuration determination and description
2p4 indicates that there are four electrons in the second (p) subshell of the second principal energy level
Read the periodic table to determine electron configuration
-lowest s is 1s, lowest p is 2p, lowest d is 3d, lowest f is 4f
Hund’s rule and implications, special elements
finding a seat on a crowded bus, electrons find their own orbital
half-filled and fully filled orbitals have more stability
chromium and copper groups are therefore exceptions to electron configuration, moving an electron from s to d
chromium = 4s13d5
copper = 4s13d10
paramagnetic vs diamagnetic
paramagnetic materials have unpaired electrons and are weakly attracted to the magnetic field
diamagnetic materials have only paired electrons and will be slightly repelled to the magnetic field
A elements and B elements
A elements are representative elements and include groups 1A through 8A (everything but transition elements and bottom of periodic table)
B elements are nonrepresentative elements and include the transition elements and lanthanide and actinide series
Effective nuclear charge trend and equation
indicates the electrostatic attraction between the valence shell electrons and the nucleus
increases from right to left, as one moves down a group principal quantum number increases and Zeff is more or less constant
Zeff = Z(atomic number) - S(non-valence electrons)
Atomic and ionic radii definition and trend
atomic radius decreases from left to right and from bottom to top
ionic radii of metals near the metalloid line is dramatically smaller than that of other metals
Ionization energy definition and trend
energy required to remove an electron from a gaseous species
-removing an electron is an endothermic process
increases from left to right and from bottom to top
groups 1 and 2 are called active metals for their low ionization energy
Electron affinity definition and trend
the energy dissipated by a gaseous species when it gains an electron, opposite of ionization energy
increases from left to right and from bottom to top
noble gases have extremely small electron affinities however
Electronegativity definition and trend
the attractive force generated in a chemical bond
increases from left to right and from bottom to top
Alkali metals
largest atomic radii, react readily with nonmetals to lose an electron
Alkaline earth metals
two electrons in valence shell
Chalcogens
Oxygen group not as reactive as halogens but crucial in biology
Halogens
desperate to complete their octets
Noble gases
inert
Transition metals
low electron affinities, ionization energies, and electronegativities have different possible oxidation states
Exceptions to the octet rule and examples
Incomplete octet hydrogen, helium, and lithium (2), beryllium (4), boron (6)
Expanded octet
-Any element in period 3 and greater can hold more than 8 electrons, including phosphorus (10), sulfur (12), chlorine (14), and others
Odd numbers of electrons Ex: NO has eleven valence electrons
Coordinate covalent bond
If both of the shared electrons are contributed by only one of the two atoms, that is a coordinate covalent bond
once it is formed it is indistinguishable from any other covalent bond
Ionic bonds in solid state
In solid state, the ionic constituents of the compound form a crystalline lattice of repeating positive and negative ions
Formal charge calculation
Valence shell - dots - dashes
Lewis structures steps
Draw out backbone with the least electronegative atom in the center
Count all the valence electrons
Complete the octets of all atoms bonded to the central atom, using the remaining valence electrons left to be assigned
Place any extra electrons on the central atom
VSEPR Theory Arrange the electron pairs around the central atom so that they are as far apart as possible
Electronic geometry
Spatial arrangement of all pairs of electrons
Molecular geometry
spatial arrangement of only the bonding pairs of electrons
determined by coordinate number -number of atoms that surround and are bonded to a central atom
Hydrogen bonds
Even hydrogen bonds have only about 10 percent the strength of a covalent bond
Nitrogen, Oxygen, or fluorine bonded to hydrogen
Moles calculation
Moles = mass of sample/ molar mass
Gram equivalent weight
the amount of a compound, measured in grams, that produces one equivalent of the particle of interest
Gram equivalent weight = molar mass/n where n is the number of particles of interest produced or consumed
Ex: gram equivalent weight of H2Co3 is half of its molar mass with interest towards h+ ions
Useful for acid-base chemistry
Normality
equivalents/L (where molarity is moles per liter)
most commonly used for hydrogen ions concentration
Ex: A 1N solution of acid would be like HCL, and 2N would be like H2So4
Molarity
Normality/n or moles per liter
Empirical formula
gives the simplest whole-number ratio
Combustion reaction
involves a fuel (usually hydrocarbon) and a oxidant (normally oxygen), forming carbon dioxide and water
Neutralization reactions
a specific type of double-displacement reaction: acid + base = water + salt
Cations and Ions naming (metals, less charge, more charge, monatomic, less oxygen, more oxygen)
For metals the charge is indicated by a Roman numeral in parentheses
- ous: less charge, -ic: greater charge
- ide: monatomic anions
Hypo- indicates less oxygen, per- indicates more oxygen
Formula and Charge:
Acetate, Cyanide Permanagante, Chromate, Dichromate, Borate, Ammonium, Thiocyanate
When is a solute considered a strong electrolyte?
A solute is considered a strong electrolyte if it dissociates completely into its constituent ions
Arrhenius equation takeaways
k = Ae^(-Ea/RT) k is the rate constant, A is the frequency factor, Ea is the activation energy of the reaction, R is the ideal gas constant, and T is the temperature in kelvins
Transition state energy
Transition state/activation complex has greater energy than both the reactants and the products and is denoted by the symbol ‡
Homogenous catalysis
the catalyst is in the same phase as the reactants
heterogeneous catalysis
the catalyst is in a distinct phase
Determination of rate law
For the general reaction aA + bB -> cC + dD, rate = k[A]^x[B]^y The values of x and y are almost never the same as the stoichiometric coefficients, the orders of a reaction must be determined experimentally
Mixed-order/broken-order reactions
refer to either non-integer orders (fractions) or to reactions with varying rate orders
fractions are specifically described as broken-order
law of mass action (determining Keq from concentration)
For a generic reversible reaction aA + bB ⇔ cC + dD, if the system is at equilibrium constant temperature
Keq = [C]^c[D]^d / [A]^a[B]^b
products over reactants
Keq = (x)^2/1-x
- if x amount of A has reacted and x amount of B and C have been produced, and 1 is the starting concentration
- can be rounded so the denominator is simply the starting concentration (in this case 1)
Types of systems and energy/matter exchange
Isolated System cannot exchange energy or matter with surroundings
Closed System can exchange energy but not matter with surroundings
Open System can exchange energy or matter with surrounds
First law of thermodynamics
Change in internal energy can only occur through heat or work
ΔU = Q - W
Isothermal processes
Constant temperature, Internal energy is constant
ΔU = 0
Adiabatic processes
no heat exchanged between the system and environment
ΔU = - W
Isobaric processes
constant pressure does not alter the first law, but appears as a flat line on a P-V graph
Isovolumetric (Isochoric) processes
constant volume
ΔU = Q
State functions
describe the system in an equilibrium state without respect to process
Pressure, density, temperature, volume, enthalpy, internal energy, Gibbs free energy, entropy
Standard conditions vs standard temperature and pressure
Standard conditions: 298K 1atm and 1 M concentrations
Standard temperature and pressure (STP): 273K and 1 atm
Phase diagrams- critical point
the temperature and pressure above which there is no distinction between the phases
-supercritical fluid
Enthalpy, Change in Enthalpy equation
equivalent to heat under constant pressure (an assumption the MCAT usually makes)
ΔHrxn = Hproducts - Hreactants
Equation for heat change, specific heat of water
q = mcΔT
specific heat of water = 1cal/g*K
definition of heat capacity
the product mc, mass times specific heat
Bomb calorimeter
constant-volume calorimetry
Because W = PΔV, no work is done in an isovolumetric process, and (ΔU = Q)
Also an adiabatic process, no heat is exchanged between the calorimeter and the rest of the universe, but it is exchanged between the steel decomposition vessel and the surrounding water
Equation for heat required for phase change
q=mL
m is mass and L is latent heat
Entropy, second law of thermodynamics
time’s arrow, entropy always increases if not hindered from doing so
Entropy equation (heat and temperature)
ΔS = Qrev/T
change in entropy = heat gained or lost in a reversible process/ Temperature in Kelvin
Gibbs free energy equation
ΔG = ΔH - TΔS get higher test scores
Free energy change equations with K
ΔGrxn = -RTlnK deriving the standard free energy change for a reaction
ΔGrxn = -RTlnQ/K deriving the free energy change for a reaction not at equilibrium K is equilibrium constant
atm to mmHG to torr to kPA
1 atm = 760mmHG = 760 torr = 100kPA
when to use STP or standard state conditions
STP is generally used for gas law calculations; standard state conditions are used when measuring standard enthalpy, entropy, free energy changes, and electrochemical cell voltage
ideal gas law
PV=nRT
density equation, density gas equation
p (density) = mass/Volume = PM/RT
density = pressure*molar mass / R*Temperature
R = 0.0821 liter·atm/mol·K
R = 8.3145 J/mol·K
constant relationships of gas exchange
PV/T is constant, PV is constant, V/T is constant, n/V is constant and equals k (moles/volume is constant)
-these can all be derived from the PV=nRT equation
Molar mass of gas calculation
M = (Density@STP) * 22.4L/mol
Molar Mass can be calculated as the product of the gases density at STP and the STP volume of one mole of gas
Partial Pressures equations
Pt = Pa + Pb + Pc
Partial pressure of gas is related to its mole fraction
Partial pressure = moles of gas A / total moles of gas
The relationship between concentration and pressure is constant
Kinetic molecular theory assumptions
explains the behavior of gases
Assumptions
1 Particles have negligible volume
2 There is no intermolecular attractions or repulsions
3 Particles are in continuous, random motion, undergoing collisions with other particles and the container walls
4 Collisions between any two gas particles are elastic, with conservation of both momentum and kinetic energy
5 The average kinetic energy of gas particles is proportional to the absolute temperature of the gas, and is the same for all gases at a given temperature, irrespective of chemical identity or atomic mass
two Average molecular speed of gas equation
KE = 1/2mv^2 = 3/2kBT proportional to 3/2 the absolute temperature of the gas and Boltzmann constant
uRMS = sqr(3RT/M) M is molar mass
Diffusion calculation from molar mass
Graham’s law: under isothermal and isobaric conditions, the rate at which two gases diffuse are inversely proportional to the square roots of their molar masses
r1/r2 = sqr(M2/M1) r is the diffusion rate, M is the molar mass
Effusion calculation
the rates of effusion are proportional to the average speeds
van der Waals equation of state purpose
corrects the ideal gas law of intermolecular attractions and molecular volume
Solvation
Solvation is the electrostatic interaction between solute and solvent molecules.
This is also known as dissolution, and when water is the solvent it can be called hydration
Sparingly soluble salts dissolve minimally in the solvent
how is H+ found in solution?
H+ is never found alone in solution, it is found bonded to an electron pair donor in a coordinate covalent bond
7 General solubility rules
All salts containing Ammonium and alkali metal cations are water-soluble
All salts containing nitrate (NO3-) and acetate (CH3COO-) anions are water-soluble
Halides excluding fluorides are water-soluble, with the exception of those formed with Ag+, Pb2+, and Hg2+
All salts of the sulfate ion (SO4) are water soluble, with the exceptions of those formed with Ca, Sr, Ba, and Pb
All metal oxides are insoluble, with the exception of those formed with the alkali metals, ammonium, and CaO, SrO, and BaO
All hydroxides are insoluble, with the exception of those formed with the alkali metals, ammonium, and Ca2+, Sr2+, and Ba2+
All carbonates, phosphates, sulfides, and sulfites are insoluble, with the exception of those formed with the alkali metals and ammonium
WE WON’T BE EXPECTED TO MEMORIZE THESE RULES, BUT ALL GROUP 1 METALS SALTS AND ALL NITRATE SALTS ARE SOLUBLE
Complex ion formation
Complexes are held together with coordinate covalent bonds
In some complexes, the central cation can be bonded to the same ligand in multiple places (chelation)
-generally requires large organic ligands that can double back to form a second (or even third) bond with the central cation
Molality
Moles/kilogram
Normality
reaction dependent, equal to the number of equivalents of interest per liter of solution
Solubility product constants and ion product
Ksp = [A^n+]^m[B^m-]^n
-does not take into account solid compounds
ion product (IP)
- analogous to the reaction quotient Q for other chemical reactions
- same form as the equation for the solubility product constant but the values aren’t at equilibrium
Common ion effect
reduction in molar solubility with the addition of common ions
Colligative properties, name them
Properties dependent on concentration but not chemical identity
Vapor pressure, boiling point elevation and freezing point depression, osmotic pressure
Vapor pressure equation
Mole fraction of solvents * Vapor pressure of solvent in pure state
Boiling point elevation and freezing point depression equation
iKm
m is molality, moles/kg
value is negative for freezing point
Osmotic pressure equation
iMRT
M is molarity, i = van’t Hoff factor
Arrhenius acid/base
dissociate H+ in solution or dissociate OH- in solution c
ontain either H at the beginning of their formula or OH at the end
Bronsted-Lowry acid/base
donates H+ or accepts H+ not limited to aqueous solutions conjugate base pairs
Lewis acid/base
electron pair acceptor (acid) or electron pair donor (base)
same chemistry as coordinate covalent bond formation, complex ion formation, or nucleophile-electrophile interactions
(nucleophile is a base, electrophile is an acid)
nomenclature anion to acid (with respect to oxygen)
if the anion ends in -ite (less oxygen), then the acid will end with -ous acid
If the anion ends in -ate (more oxygen), then the acid will end with -ic acid
Water dissociation constant Kw
Kw = [H3O+][OH-] = 10^-14 @ 298K
Estimating pH
if H+ is .001 or 10^-3 , pH = 3 -log(n*10^-m) = m-1.10-n or m-.n dissociation of strong acids and bases is said to go to completion
Strong acids
HCl (Hydrochloric acid)
HBr (Hydrobromic acid
H2SO4 (Sulfuric acid)
HNO3 (Nitric acid)
HCLO4 (Perchloric Acid)
Strong bases
NaOH (sodium hydroxide)
KOH (potassium hydroxide)
Activity of Weak Acids and Bases (Ka and Kb)
Ka = [H3O][A-] / [HA]
Kb = [B+][OH-] / [BOH]
Ka and Kb are mostly used to determine the concentration of one of the species at equilibrium
Salt formation, varying strengths acids + varying strengths bases
Strong acid + strong base: HCl + NaOH -> NaCl + H2O
Strong acid + weak base: HCl + NH3 -> NH4Cl
Weak acid + strong base: HCLO + NaOH -> NaClO + H2O
Weak acid + weak base: HCLO + NH3 -> NH4ClO
Understanding titrations
At the equivalence point, the number of equivalents of acid and base are equal
NormalityA * VolumeA = NormalityB * VolumeB
Indicators are weak organic acids or bases that have different colors in their protonated and deprotonated states
Polyvalent acids and bases yield multiple equivalence points
-at each half-equivalence points, half of a given species has been protonated or deprotonated
Bicarbonate buffer system
Co2 <=> H2CO3 ⇔ H+ + HCO3-
Henderson Hasselbach equation
For a weak acid buffer solution
pH = pKa + log [A-]/[HA]
For a weak base buffer solution
pOH = pKb + log [B+]/[BOH]
remember concentration ratio as products over reactants
If the concentration of both the acid and conjugate base were double the buffering capacity would double, not the pH
Assigning oxidation numbers
the oxidation number of a free element is zero
the oxidation number for a monatomic ion is equal to the charge of the ion
the sum of the oxidation numbers in a compound must add to that compounds charge
Electrochemical cells, location of oxidation/reduction and pathway of electron
contained systems in which oxidation-reduction reactions occur
oxidation always occurs at the anode, and reduction occurs at the cathode
electron always travels from anode to cathode
Galvanic cell, description, charge of anode and cathode, spontaneity etc
all the batteries you own
spontaneous, + emf, -ΔG
in a galvanic cell the anode is negative and the cathode is positive
has an aqueous electrolyte solution composed of cations and anions, the charge gradient is dissipated by the presence of a salt bridge, which contains ions that will not react with the electrodes or with the ions in solution
Electrolytic cell, description, charge of anode and cathode, spontaneity etc
opposite of galvanic cells, except for An Ox Red Cat and anode-> cathode
driven by an external battery source, known as electrolysis
nonspontaneous, - emf, +ΔG
in an electrolytic cell the anode is positive and the cathode is negative
Construction of cell diagram
shorthand notation representing the reactions in an electrochemical cell
Ex: Zn (s) | Zn2+ (1M) || Cu2+ (1M) | Cu (s)
- The reactants and products are always listed from left to right in anode | anode solution (concentration) || cathode solution (concentration) | cathode
- A single vertical line indicates a phase boundary whereas a double vertical line indicates the presence of a salt bridge or other barrier
Faraday constant
10^5 C/mol e-
Electrodeposition equation
mol M = It/nF
mol metal ion being deposited = current * time / 10^5 * number of electron equivalents for a specific metal ion
OR
mol e- = It/10^5
Concentration cells
a type of galvanic cell, contains two half-cells connected by a conductive material allowing a spontaneous oxidation-reduction to proceed, which generates a current and delivers energy
Electromotive force
Describes the electrochemical cell, if the emf is positive, the cell is spontaneous
Ecell = Ered,cathode - Ered,anode
Gibbs free energy of electrochemical cell
ΔG* = -nFE*cell where ΔG* is the standard change in free energy, n is the number of moles of electrons exchanged, F is the Faraday constant, and E*cell is the standard emf of the cell
Nernst equation
Ecell = E*cell - (RT/nF)lnQ
Q = [C]c[D]d / [A]a[B]b
Reaction equilibrium equation
ΔG* = -RTlnKeq
Nomenclature, Aldehydes
O=C-H
indicated with the suffix -al or the prefix oxo- when not the highest priority group
Nomenclature, Ketones
R-(C=O)-R
indicated with the suffix -one
Nomenclature, Esters
O=C-OR
indicated with the suffix -oate
Nomenclature, Amides
O=C-N indicated with the suffix
-amide
Nomenclature, Anhydrides
O=C-O-C=O indicated with the suffix
-anhydride
Nomenclature, Imine
C=N
Nomenclature, Enamine
C=C-N
Nomenclature, Cyanohydrins
HO-C-Cn
Nomenclature, Aldol
O=C-C-C-OH aldehyde plus alcohol
Conformational vs configurational isomers
Conformational isomers differ in rotation, Configurational isomers require bond breaking to interconvert (enantiomers and diastereomers)
Newman projection different conformations
eclipsed
staggered
- anti - two largest groups are on opposite sides
- gauche - two largest groups are 60* apart
Enantiomers
Enantiomers differ at all chiral centers
when present in equal concentrations they form racemic mixtures, which can be separated with a single enantiomer of another compound, leading to two diastereomers
Diastereomers
Cis-Trans molecules
differ in arrangment around an immoveable bon
Cylic conformations and ring strain factors
Cyclic conformations can be either stable or unstable depending on ring strain
Angle strain bond angles deviate from their ideal values
Torsional strain cyclic molecules are eclipsed or gauche
Nonbonded strain (van der waals repulsion) nonadjacent atoms or groups compete for the same space
Meso compounds
a molecule with at least 2 chiral centers that has an internal plane of symmetry
(E) and (Z) form
polysubstituted double bonds
Z if the highest priority (atomic number) are on the same side, E if opposite
(R) and (S) forms
highest atomic number is highest priority
clockwise is R where counterclockwise is L
switch if the lowest priority is in front of molecule (wedge in Fischer, side in skeleton)
bond strength/acidity and periodic table
acidity increases towards the bottom right of the periodic table
Nucleophiles
either lone pairs or pi bonds that can form new bonds to electrophiles
determined by four major factors
Four factors of nucleophilicity
Charge - nucleophilicity increases with electron density / negative charge
Electronegativity - nucleophilicity decreases as electronegativity increase because these atoms are less likely to share electron density
Steric hindrance - bulkier molecules are less nucleophilic
Solvent - protic solvents can hinder nucleophilicity by protonation of the nucleophile or through hydrogen bonding
For protic solvents, nucleophilicity goes I- > Br- > Cl- > F-
For aprotic solvents, nucleophilicity goes I- < Br- < Cl- < F-
Nucleophilic Substitutions SN1 description and rate determination
Unimolecular, contain two steps
rate of the reaction depends only on the concentration of the substrate rate = k[R-L], where R-L is an alkyl group containing a leaving group
Nucleophilic Substitutions SN2 description and rate determination and special considerations
Bimolecular, contain only one step
rate of the reaction depends on both the concentration of the substrate and the nucleophile
rate = k[R-L][Nu:]
nucleophile must be strong and the substrate cannot be sterically hindered
accompanied by an inversion of relative configuration, the position of the substituents around the substrate carbon will be inverted (R -> S or vice versa)
Where will a redox reagent preferably act
A redox reagent will tend to act on the highest priority functional group
Phenol nomenclature
adjacent carbons are called ortho, separated by a carbon is called meta, opposite sides is called para
Alcohol oxidation reactions (Primary and secondary)
Primary alcohols can be oxidized to aldehydes
-only be PCC, a mild oxidant
Primary alcohols can be oxidized past aldehydes to geminal (same carbon) diols to carboxylic acid
-with stronger oxidizing agents than PCC
Secondary alcohols can be oxidized to ketones with any oxidizing agent
Tertiary alcohols cannot be oxidized
Mesylates and tosylates
Mesylate have the functional group -SO3CH3
Tosylates contain the functional group -SO3-C6H4(benzene)-CH3
these groups can serve as protecting groups or can be formed from alcohols to make better leaving groups
Quinones and hydroxyquinones
Treatment of phenols with oxidizing agents produces quinones (O=ring) example of secondary alcohol oxidation
not necessarily aromatic because they lack the conjugated ring structure
Hydroxyquinones share the same ring and carbonyl backbone as well as a hydroxyl group
Ubiquinone (CoQ) is a biologically active quinone
Aldehydes and Ketones- Hydration
In the presence of water, aldehydes and ketones react to form geminal diols (same carbon diols)
-the reaction rate can be increased by adding a small amount of catalytic acid or base
Acetal or ketal formation
Acetals C-(OR-)C(-OR)-H or ketals C-(OR-)C(-OR)-C can be created with the reaction of aldehydes or ketones with two equivalents of alcohol or a diol catalyzed by anhydrous acid
Hemiacetal or hemiketal formation
Hemiacetals C-(OR-)C(-OH)-H or hemiketals C-(OR-)C(-OH)-C can be created with the reaction of aldehydes or ketones with one equivalent of alcohol
endpoint in basic conditions
Imines and enamines formation
Simplest case of imine formation from aldehyde/ketone:
-ammonia adds to the carbon and water is lost
–examples of a condensation reaction and nucleophilic substitution
Imines can tautomerize to enamines
Cyanohydrin formation
HCN + aldehydes or ketones can create cyanohydrins, as the CN group adds to the carbonyl carbon and forces reduction
Oxidation of aldehydes
Aldehydes can be oxidized to carboxylic acids with any oxidizing agent stronger than PCC
Aldehyde and ketone reduction by hydride reagents
Aldehydes and ketones can be reduced to form alcohols
-most often performed with the hydride reagents LiALH4 and NaBH4
Keto-Enol tautomerization
the enol is named for its C=C and -OH the two isomers are called tautomers, and interconvert through tautomerization/enolization
the thermodynamic intermediate is most stably substituted but forms less quickly
Enamines are tautomers of imines in the same way
Aldol condensation
An aldehyde or ketone acts as both an electrophile (in keto form) and a nucleophile (in enolate form), and the end result is a formation of a carbon-carbon bond from a beta carbon to an alpha carbon
-requires a catalytic amount of base
Aldol has a carbonyl and a hydroxyl group sharing a beta carbon
-aldehyde and alcohol = aldol
An example of a condensation reaction and a dehydration reaction because two molecules are joined with the loss of a small molecule
Retro-aldol reaction
reverse of aldol condensation, aqueous base (OH-) is added and heat is applied
Carboxylic Acid varying acidity
Groups like -NO2 or halides are electron withdrawing and increase acidity
-Dicarboxylic acids are therefore more acidic than analogous monocarboxylic acids due
–beta-dicarboxylic acids have high acidity of the alpha-carbon that they share, although not as high as the hydroxyl hydrogens
Groups like -NH2 or -OCH3 are electron donating and decrease the acidity
Carboxyl Acid Synthesis
As described above, oxidation of aldehydes and primary alcohols with strong oxidizing agents ((N)Cr2O7, CrO3, KMnO4)
Nucleophilic acyl substitution of carboxylic acid
Nucleophile- can add to carboxylic acid and substitute for -OH while donating a proton
Different than aldehydes or ketones because those participate in addition
Weak bases (Conjugate bases of strong acids) make good leaving groups
carboxylic acid + amine
Amides
- Carboxylic acids can be converted into amides with the addition of an amine in an acidic or basic solution
- Amides that are cyclic are called lactams and replace -oic with -lactam
carboxylic acid + alcohol; nomeclature if product is cyclic
A reaction of carboxylic acids with alcohol under acidic conditions results in formation of an ester, with the R- of the ester replacing the -H of the carboxylic acid’s hydroxyl group
- occurs most rapidly with primary alcohols
- in acidic conditions
Condensation reaction with water side product
Esters that are cyclic are called lactones and replace -oic with -lactone
carboxylic acid + carboxylic acid
Anhydrides can be formed by the condensation of two carboxylic acids
Carboxylic acid reduction
Carboxylic acids can be reduced to primary alcohols by the use of LiALH4
Decarboxylation
beta-keto acids like 1,3-dicarboxylic acids may spontaneously decarboxylate when heated
Saponification
Long chain carboxylic acids reacted with sodium or potassium hydroxide to create a salt (H of the hydroxyl is replaced with Na or K)
Know the product as soap
Carboxylic acid derivates
Amides, Esters and Anhydrides are all formed by a condensation reaction with a carboxylic acid
Amide synthesis
synthesized by reaction of other carboxylic acid derivatives with either ammonia or an amine
Ester synthesis
dehydration synthesis products of other carboxylic acid derivatives and alcohols
Fischer esterification
Under acidic conditions, mixtures of carboxylic acids and alcohols will condense into esters with the R replacing the H of the hydroxyl group
Anhydride synthesis
condensation dimers of carboxylic acids, wherein the oxygen of one carboxylic acids hydroxyl group bonds to the alpha carbon of another
Relative reactivity of carboxylic acid derivatives
Anhydrides are most reactive, followed by esters (tied with carboxylic acids although they usually have lower boiling points) and then amides
Strain in cyclic derivatives of carboxylic acids
certain times lactams and lactones are more reactive to hydrolysis because they contain more strain
four membered rings have torsional strain from eclipsing interactions and angle strain from not having a 109.5* angle
anhydride cleavage
Addition of a nucleophile to an anhydride results in carboxylic acid derivative and carboxylic acid
-derivative depends on nucleophile; if amine the derivative is an amide
transesterification
Alcohols can act as nucleophiles and displace the esterifying group on an ester, transforming one alcohol for another
Amide hydrolysis
Amides can be hydrolyzed under highly acid conditions via nucleophilic substitution, replacing an -NH2 with a -OH from water
Strecker synthesis
generates an amino acid from an aldehyde
An aldehyde is mixed with ammonium chloride (NH4Cl) and potassium cyanide.
The ammonia attacks the carbonyl carbon, generating an imine.
The imine is attacked by the cyanide, generating an aminonitrile
The aminonitrile is hydrolyzed by two equivalents of water, generating an amino acid
Gabriel synthesis
generates an amino acid from potassium phthalimide, diethyl bromomalonate, and an alkyl halide
phthalimide attacks the diethyl bromomalonate, generating a phthalimid-omalonic ester
The phtalimidiomalongic ester attacks an alkyl halide, adding an alkyl group to the ester
The product is hydrolyzed, creating phthalic acid and converting the esters into carboxylic acids
One carboxylic acid of the resulting 1,3-dicarbonyl is removed by decarboxylation
IR spectroscopy; fingerprint area, hydroxyl area, carboxylic acid area, aldehyde/ketone area, amine
fingerprint region of 1500 to 400 cm-1
Hydroxyl group at 3300 for alcohols, 3000 for carboxylic acids with broad peak
carbonyl at 1700 with sharp deep peak (aldehydes and ketones)
NH bonds at 3300 with sharp peak
UV spectroscopy
never have to interpret but is obtained by passing light through a sample and recording the absorbance
NMR spectroscopy, which direction is downfield what affect do electronegative atoms have what is height and number of peaks proportional to
left is downfield and increasing chemical shift TMS is marked at 0, ignore it
The height of each peak is proportional to the number of protons it contains
electronegative atoms pull electron density downfield
the number of peaks indicates the number of adjacent protons
NMR spectroscopy, aldehyde hydrogens, carboxylic acid hydrogen, aromatic hydrogen
Aldehyde Hydrogens > 9 to 10 ppm
Carboxylic acids Hydrogens > 10.5 to 12 ppm
Aromatics Hydrogens > 6.0 to 8.5 ppm
NMR spectroscopy, sp3 hydrogens, sp2 hydrogens, sp hydrogens
sp3 hydrogens -> 0 to 3 ppm
sp2 (alkene) hydrogens -> 5 to 7 ppm
sp hydrogens -> 2 to 3 ppm
Extractions
have an aqueous phase and a organic phase
once the desired product has been isolated in the solvent, we can obtain the product alone by evaporating the solve, usually with a rotovap
Filtrations, gravity vs vacuum
isolates a solid (residue) from a liquid (filtrate)
Gravity filtration is used when the product of interest is in the filtrate, hot solvent is used to maintain solubility
Vacuum filtration is used when the product of interest is solid. A vacuum is connected to the flask to pull the solvent through more quickly
Distillations, simple vs vacuum vs fractional
separates liquids according to their boiling point
Simple distillation if the boiling points are under 150C and are at least 25C apart
Vacuum distillation if the boiling points are over 150C to prevent degradation of the product
Fractional distillation if the boiling points are less than 25C apart because it allows more refined separation of liquids by boiling point
Chromatography general principle
Use two phases to separate compounds based on physical or chemical properties
the stationary phase or adsorbent is a polar solid
the mobile phase runs through the stationary phase and is usually a liquid or gas
TLC and paper chromatography stationary phase and mobile phase
TLC and paper chromatography the stationary phase is a polar material, such as silica, alumina, or paper the mobile phase is a nonpolar solvent, which climbs the card through capillary action
Reverse phase chromatography
need to insert
Rf
Rf- Retardation factor distance spot moved / distance solvent front moved
Column chromatography general principle
uses an entire column filled with silica or aluminum beads as an adsorbent
ion exchange chromatography, size exclusion chromatography, affinity chromatography
Ion exchange, size exclusion, and affinity chromatography
Ion exchange chromatography
the beads are coated with charged substance to bind compounds with opposite charge
Size exclusion chromatography
the beads have small pores which trap smaller compounds and allow larger compounds to travel through faster
Affinity chromatography
the column is made to have high affinity for a compound by coating the beads with a receptor or antibody
Gas chromatography
separates vaporizable compounds according to how well they adhere to the adsorbent in the column
the stationary phase is a coil of crushed metal or a polymer, the mobile phase is an inert gas
can be combined with mass spectroscopy to ionize the fragments and determine molecular weight or structure
HPLC (High performance liquid chromatography)
similar to column chromatography but uses sophisticated computer-mediated solvent and temperature gradients
