General Chemistry Reference Sheet
This reference sheet addresses some of the more peculiar
pieces of information that need to be memorized in a general chemistry course. It also contains a simple set of essential formulas in chemistry with cautions, explanations,
and general tips.
This sheet is meant to be as concise as possible, and many
information in the textbook is left out in favor of cautions
and tips. This sheet is, therefore, best used as a supplement to, not a replacement of, the textbook.
Soluble
NO3 −
CH3 COO−
Cl−
Br−
I−
SO4 2−
Insoluble
S2−
2−
SI Fundamental Units
Mass
Length
Time
Temperature
Amount of substance
Electric current
Luminous intensity
Kilogram (kg)
Meter (m)
Second (s)
Kelvin (K)
Mole (mol)
Ampere (A)
Candela (cd)
Atomic Experiments and Models
Discovered e− ; Cathode ray
Plum pudding model
R. A. Millikan
Measured charge of e− ; Oil drop
H. Becquerel/M. Curie
Discovered radioactivity
E. Rutherford
Discovered α, β, and γ rays
Discovered nucleus; Gold foil experiment
J. Chadwick
Discovered neutrons
N. Bohr
Bohr model (electron orbits)
Quantum mechanicists
Quantum model
J. J. Thomson
Polyatomic Ions
+
−
CO3
PO4 3−
OH−
Exceptions
None
None
Ag+ , Hg2 2+ , Pb2+
Ag+ , Hg2 2+ , Pb2+
Ag+ , Hg2 2+ , Pb2+
Sr2+ , Ba2+ , Hg2 2+ , Pb2+
Exceptions
NH4 + , alkali metal cations, Ca2+ ,
Sr2+ , Ba2+
NH4 + , alkali metal cations
NH4 + , alkali metal cations
NH4 + , alkali metal cations, Ca2+ ,
Sr2+ , Ba2+
Strong acids and bases dissociate in water completely.
Strong Acids
HCl
HBr
HI
HClO3
HClO4
HNO3
H2 SO4
Ba(OH)2
Strong Bases
Alkali metal hydroxides
Ca(OH)2
Sr(OH)2
Activity Series
Metals below H+ cannot react with acids to form H2 . More
active metals are better reducing agents.
From most active to least active:
Li+ , K+ , Ba2+ , Ca2+ , Na+ , Mg2+ , Al3+ , Mn2+ , Zn2+ ,
Cr3+ , Fe2+ , Co2+ , Ni2+ , Sn2+ , Pb2+ , H+ , Cu2+ , Ag+ ,
Hg2+ , Pt2+ , Au3+
OH
C2 O4 2−
CNO−
C2 H3 O2 −
hydroxide
oxalate
cyanate
acetate
SCN−
SO3 2−
CO3 2−
PO4 3−
S2 O3 2−
NO3 −
ClO4 −
ClO3 −
ClO−
HPO4 2−
CrO4 2−
Cr2 O7 2−
MnO4 −
N3 −
C4 H4 O6 2−
O2 −
PO2 3−
SiO3 2−
chromate
dichromate
permanganate
azide
tartrate
superoxide
hypophosphite
silicate
H3 O+
PO3 3−
Hg2 2−
C2 2−
S2 2−
AsO3 3−
AsO4 3−
P2 O7 4−
nitrate
perchlorate
chlorate
hypochlorite
hydrogen
phosphate
hydronium
phosphite
mercury (I)
carbide
Phase Changes
disulfide
From solid
arsenite
To solid
arsenate
To liquid melting
pyrophosphate To gas
sublimation
Flame Colors
Calcium
Copper (I)
Copper (II)
Potassium
Lithium
Sodium
Strontium
Barium
Iron (III)
Cesium
Indium
Lead
Rubidium
Copper (II)
Nickel
Permanganate
Chromate
Dichromate
Iron (II)
Iron (III)
Blue
Green
Purple
Yellow
Orange
Light blue
Rusty yellow
Thermodynamic Laws
First Law: Energy cannot be created nor destroyed. It
can only be transferred in the form of either heat or work.
Second Law: Any spontaneous reaction increases the entropy of the universe.
Third Law: An ideal solid crystal at 0 K has an entropy
of 0.
Thermodynamic Formulas
Strong Acids and Bases
ammonium
cyanide
peroxide
hydrogen sulfate
thiocyanate
sulfite
carbonate
phosphate
thiosulfate
NH4
CN−
O2 2−
HSO4 −
Solution Colors
Ionic Solubility Chart
Brick red
Blue
Green or blue-green
Lilac
Dark red
Bright yellow
Red
Light green
Gold
Blue–Violet
Blue
Blue
Red–Violet
Standard thermodynamic conditions 298 K; 1 atm; 1 M
Kinetic energy
K = mv 2 /2
Electrostatic potential energy
UE = (kC Q1 Q2 )/d
Internal energy
∆E = q + w
Enthalpy
H = E + PV
Specific heat
s = q/(m · ∆T )
Entropy in reversible reaction
∆Ssystem = (∆H)/T
∆Ssurrounding = −(∆H)/T
Microstate-entropy relationship
S = k ln W
Gibbs free energy
G = H − TS
Gibbs free energy change
∆G = ∆H − T ∆S
∆G = ∆G◦ + RT ln Q
Hess’s Law
∆Htotal = Σ∆Hi
Constants
Boltzmann’s constant kB = 1.381 × 10−23 m2 kg · s−2 K−1
Coulomb’s constant kC = 1/(4π0 ) = 8.988 × 109 J · m/C2
Avogadro’s number
NA = 6.022 × 1023 mol−1
Faraday’s constant
F = 9.649 × 104 C/mol
Planck’s constant
h = 6.626 × 10−34 J·s
Ideal gas constants
R = 0.0821 (L · atm)/(mol · K)
R = 8.314 J/(mol · K)
Vacuum permittivity 0 = 1/(µ0 c2 ) = 8.854 × 10−12 F/m
Vacuum permeability
µ0 = 1.257 × 10−6 N·A−2
Atomic mass
1 amu = 1.661 × 10−24 g
Electron charge
e = 1.602 × 10−19 C
Electronvolt
1 eV = 1.602 × 10−19 J
Atmospheric pressure
1 atm = 1.013 × 105 Pa
Absolute zero
0 K = -273.15 ◦ C
Speed of light in vacuum
c = 2.998 × 108 m/s
Quantum Mechanical Formulas
From liquid
freezing
vaporization
From gas
deposition
condensation
-
Energy of a quantum
E = hν
Wavelength-frequency relationship RRR
c=ν·λ
Probability distribution
PV =
|ψ(x, y, z)|2 dxdydz
V
Laws of Quantum Mechanics
Heisenberg’s Uncertainty Principle
Pauli Exclusion Principle
Hund’s Rule
Atomic size
Corollary: It is impossible to determine both the position and
the momentum for a sufficiently
small particle like an e− .
No two e− in an atom can share
the same set of four quantum
numbers.
Corollary: A suborbital can hold
a maximum of 2e− .
Energy is the lowest when the
number of e− with the same spin
is maximized.
Corollary: e− will first half-fill
all the empty suborbitals, then
go back and fill the half suborbitals.
Ionic size
−
Quantum Numbers of e
Principal (n)
Azimuthal (l)
Magnetic (ml )
Spin (ms )
The energy shell of the e− , e.g. 4 in
4d1 .
The suborbital shape, with s=0,
p=1, d=2, f=3, e.g. 2 in 4d1 .
The suborbital, ranging from −l to
l, e.g. −2 in 4d1 .
The spin of e− . Two e− in the same
suborbital has either −1/2 or 1/2.
Molecular Geometry
Hybridization
sp
Atomic Properties
∆x · ∆p ≥ h̄/2
Nonbonding
electrons
Geometry
IMFs in Molecular Solids
1
the
2
N-th ionization energy
Electron affinity
Metallic character
distance between two adjacent nucleii.
Cations are smaller than their
parent atoms.
Anions are
larger than their parent atoms.
The energy required to remove the n-th electron from a
ground state gaseous atom.
The energy released by adding
an electron to a gaseous atom.
The qualities of a metal.
Metals are shiny and heat and
electricity-conducting;
they
have malleble solid form, form
basic ionic oxides, and tend
to form cations in an aqueous
solution.
Periodic Properties
Property
Atomic size
Ionization energy
Electron affinity
Metallic character
Left to Right
Decreasing
Increasing
Large if adding
to a previously
empty orbital
Decreasing
Type
IMFs
Molecular Van der Waals forces,
dipole-dipole interactions,
hydrogen
bonds
Covalent- Covalent bonds
network
linear
sp
0
1
trigonal planar
bent
Ionic
Electrostatic interactions
sp3
0
1
2
tetrahedral
trigonal pyramidal
bent
Metallic
Metallic bonds
0
1
2
3
trigonal bipyramidal
seesaw
T-shaped
linear
0
1
2
octahedral
square pyramidal
square planar
3
sp d
sp3 d2
Increasing
Types of Crystalline Solids
0
2
Top to Bottom
Increasing
Decreasing
No
apparent
change
Properties
Soft, low melting
point, poor conduction
Very hard, high
melting point, poor
conduction
Hard, high melting
point, poor conduction
Soft to very hard,
low to very high
melting point, excellent conduction
Boiling Points of Molecular Compounds
The relative boiling points of molecular compounds can be
determined by their IMFs. The stronger the IMFs are,
the higher the boiling point. (Note that linear compounds
like straight-chain hydrocarbons have higher van der Waals
forces than non-linear compounds because their molecules
have a greater area of contact.)
2
London
dispersion
force (van der Waals
forces, induced dipoledipole interactions)
Dipole-dipole
tions
interac-
Hydrogen bonds
Interactions between dipoles
partially charged through the
movement of shared electron.
Presents in all compounds.
Weakest of the three.
Interactions between dipoles
partially charged through the
electronegativity difference of
two bonding atoms.
A special kind of dipole-dipole
interactions present in compounds that have hydrogen
and either oxygen or nitrogen.
Strongest of the three.
Acid-Base Theories
Arrhenius
BrønstedLowry
Proton
donors
Proton
acceptors
Hydrogen
atom
Acids
[H+ ] >[OH− ]
Bases
[OH− ] >[H+ ]
Acids have
H+
Bases have
OH−
Unshared
electron pair
Acid + Base
→
Salt + H2 O
Conjugate
acid
+
Conjugate
base
Lewis
Electron
acceptors
Electron
donors
Electron
accepting
atom
Unshared
electron
pair
Kinetic Molecular Theory (KMT)
There is a very large number of particles;
Particles are in constant random motion and collide constantly with the wall;
Collisions of particles with the wall are perfectly elastic;
Particles exert no force upon each other.
Properties of Solutions
Solvation The uniform dispersion of a solute in a solvent.
Hydration
Solvation in water.
Crystallization
The reverse reaction of solvation.
Saturated
A solution in equilibrium.
Unsaturated A solution with less solute than saturation.
Supersaturated
A solution with more solute than
saturation. (Will undergo crystallization
if a crystal seed is present.)
Miscible
Two liquids that dissolve in any proportion.
Henry’s Law
S ∝ P (S: solubility)
Colligative Properties
Activation Energy
Physical properties of a solution that depends on the concentration of solutes. More solutes will lead to:
Collision model
1. Lower vapor pressure: PA = XA PA◦ (Raoult’s
Law)
2. Higher boiling point: Tb = Tb◦ + kb m (Molality)
3. Lower freezing point: Tf = Tf◦ − kf m (Molality)
4. Higher osmotic pressure: π = RT · M (Molarity)
Colligative properties also depend on the van ’t Hoff factor
(i = number of particles after reaction / number of particles before reaction). The greater i is, the more colligative
properties it exerts on the solution.
Reaction Rate
The reaction rate r = d[X]/dt can be determined from the
reaction by the rate law
r = k[A]a [B]b ...
Where a, b, etc. are reaction orders for the reactants. Reaction orders can only be determined experimentally, because reactions will in theory go through several steps, the
slowest of which is the rate-determining step. Reaction order is determined by the number of atoms participating in
the rate-determining step. The sum of these orders is the
overall order.
Concentration function [X] can be determined as
Z
t
rdτ + [X]0
[X]t =
0
Therefore, for first-order reactions,
ln[X]t = −kt + ln[X]0
Graphically, t is proportional to ln[X]t with the slope −k.
And for second-order reactions,
1
1
= kt +
[X]t
[X]0
Graphically, t is proportional to 1/[X]t with the slope k.
Reaction Half-time
The half-time of a reaction is the amount of time needed
to consume half of the reactants. It is denoted t1/2 .
For first order reactions, t1/2 ≈ 0.693/k. For second order
reactions, t1/2 = 1/(k[X]0 ).
Concentration
Reactions occur as a result of collisions
between molecules.
Activation energy (Ea )
The minimum energy required
for a reaction to occur.
Arrhenius equation
ln k = ln A − Ea /RT
(This means that ln k ∝ 1/T )
Activation energy is lowered when a catalyst is present. Inorganic catalysts usually provide a site on which reactants
can adsorb; organic catalysts, or enzymes, bind specific to
substrate molecules (“lock-and-key”).
Definition
(Moles solute)/(Liters solution)
(Moles solute)/(Kilogram solvent)
(Moles solute)/(Moles solution)
(Mass of solute)/(Mass of solution)
(Volume of solute)/(Volume of solution)
Caution: It is an extremely common mistake to confuse
molarity with molality. Check your R’s and L’s!
Spectrophotometry of Concentration
An equilibrium reaction will spontaneously balance an outside effect added to it. For example,
Change in amount of reactants or products: The
reaction will consume more of the substance in excess to
balance the change;
Change in volume or pressure: The reaction will form
more gas if volume increases or if pressure decreases, and
will form less gas if volume decreases or if pressure increases;
Change in temperature: Endothermic reactions will
shift left for lower temperatures and shift right for higher
temperatures. Exothermic reactions will shift right for
lower temperatures and shift left for higher temperatures.
Beer’s law, A = lc, relates concentration and light absorption.
Absorbance (A)
− log10 (I/I0 ) (liquids)
− ln(I/I0 ) (gases)
Absorption coefficient ()
Depends on the solution.
Length of path (l)
The length of the path travelled
by light.
Concentration (c)
The concentration of the solution.
In spectrophotometry, the length of path is fixed. Therefore, when using the same solution, A ∝ c.
Gas Laws
STP
273 K; 1 atm
Boyle’s Law
P ∝ 1/V
Charles’s Law
P ∝T
Avogadro’s Law
P ∝n
Ideal Gas Equation
P V = nRT
Law of Partial Pressure
pPn = Xn Pt
(3RT )/M
Effusion Rate
u= p
Graham’s Law
u1 /u2 = M2 /M1
Density Formula
d = (P M)/(RT )
Deviation from Ideal Behavior
(P V )/(RT )
Van der Waals Equation
P + (n2 a)/V 2 (V − nb) = nRT
Clausius-Clapeyron Equation ln P = −∆Hvap /(RT ) + C
Caution: When using the effusion rate formula, the R
value must be in joules (8.314), and the M value must be
converted to kg/mol.
Equilibrium Formulas
Ion-product constant of water
Kw = [H+ ][OH− ] = 1.0 × 10−14 (278 K)
Henderson-Hasselbalch equation
pH = pKa + log([base]/[acid])
Van ’t Hoff equation
d(ln K)/dT = (∆H ◦ )/(RT 2 )
ln K = −∆H ◦ /(RT ) + ∆S ◦ /R
3
Notation
Molarity (M)
Molality (m)
Mole fraction (X)
Mass percentage
Volume percentage
Le Châtelier’s Principle
Equilibrium Constant
For reactions in a solution, the equilibrium constant of a
reaction
* sS + tT
aA + bB )
Is defined as
Kc =
[S]s [T ]t
[A]a [B]b
When the reaction is in equilibrium.
If the reaction is a equilibrium between a solid and its ions
in solution, then Kc is its solubility product constant, Ksp .
If the reaction is a dissociative reaction of a weak acid, then
Kc is its acid dissociation constant, Ka . Polyprotic (having more than one H) acids have multiple Ka , but usually
Ka1 determines the pH.
For a conjugate acid-base pair, Ka · Kb = Kw .
If the reactants are gases, then the equilibrium constant is
defined as
P sP t
Kp = Sa Tb
PA PB
When the reaction is in equilibrium.
The Formulas above, when applied to non-equilibrium situations, gives Q. The reaction forms products if Q < K,
reactants if Q > K, and nothing if Q = K (already in
equilibrium).
Equilibrium Constant (cont.)
Types of Magnetic Materials (cont.)
Stereoisomerism (cont.)
For all equilibrium reactions, there are more reactants than
products if K < 1, more products than reactants if K > 1,
and the same amount of reactants and products if K = 1.
Diamagnetic: Can be magnetized to repulse external
magnetic fields, but cannot retain magnetism. Diamagnetic materials have a magnetic permeability of less than
µ0 . Examples of diamagnets are bismuth and antimony.
Ferromagnetic: Can be magnetized and retain magnetism. Ferromagnetism depends both on the chemical
composition and the structure of the material (iron is a
ferromagnet, while stainless steel is not). Examples of ferromagnets include cobalt and iron.
E/Z isomerism: If the two ends of the bond do not have a
common hydrogen atom, then the compound exhibits E/Z
isomerism. The Z isomer has the “larger” substituents
(defined by the CIP Rules) of both ends on the same side,
while the E isomer has the larger substituents on different
sides.
Acid Character of Hydrogen Atoms
Hydrogen atoms are acidic when they are weakly bonded,
and when the molecule/atom they are bonded to forms
stable anions.
In organic compounds, the hydrogen atoms in carboxyl
groups (COOH) are usually the most acidic.
Indicators for Acid-Base Titration
Indicator
Small pH
Methyl violet
Yellow
Bromophenol
Yellow
blue
Methyl orange
Red
Methyl red
Red
Litmus
Red
Bromothymol
Yellow
blue
Phenolphthalein Colorless
Color change
0.0–1.6
3.0–4.6
Large pH
Violet
Blue
3.1–4.4
4.4-6.2
5.0–8.0
6.0–7.6
Yellow
Yellow
Yellow
Blue
8.3–10.0
Pink
Oxidation and Reduction
Oxidation
Loss of electrons
Oxidation num. increases
Occurs at anode
Mnemonic devices:
Reduction
Gains electrons
Oxidation num. decreases
Occurs at cathode
• “OIL RIG” (Oxidation Is Loss, Reduction Is Gain)
• “What an ox loses, a red cat gains” (An =
anode; ox = oxidation; red = reduction; cat = cathode)
Electrochemical Formulas
E = −(∆G)/(nF )
E = E ◦ − (RT /nF ) ln Q
E = E ◦ − (0.0592/n) log Q
◦
◦
◦
Standard cell potential
E = Ered
cathode − Ered
anode
Energy of a charged particle
E = qV
Faraday’s Law of Electrolysis
m = (Q/F )(M/z)
Electromotive force
Nernst equation
Types of Magnetic Materials
Paramagnetic: Can be magnetized to attract external
magnetic fields, but cannot retain magnetism. Paramagnetic materials have a magnetic permeability of more
than µ0 . They usually have free electrons, especially d
and f electrons. Their magnetization follows Curie’s Law
(M = C · B/T ). Examples of paramagnets are tungsten
and cesium.
Nuclear Chemistry
Alpha particles (α)
Helium nuclei (42 He)
−
Beta particles (β )
Electrons (0−1 e)
Positrons (β + )
Antielectrons (01 e)
Gamma radiation (γ)
High energy radiation (00 γ)
Units of radioactivity
SI: Becquerel (Bq): 1 nucleus/s
(Disintegration per second) Curie (Ci): 3.7 ×1010 nuclei/s
Units of absorbed radiation
SI: Gray (Gy): 1 J/kg
(Energy per kilogram tissue)
Rad: 0.01 Gy
Metallurgy
Metallurgy is the extraction of minerals from ores.
Pyrometallurgy: The use of heat to convert ores to metals. (Example: Production of iron)
Hydrometallurgy: The use of chemical processes in a solution to separate a metal from its ore. (Example: Bayer
process for producing aluminum)
Electrometallurgy: The use of electrochemical processes
to separate a metal. (Example: Hall process for producing
aluminum)
Hydrocarbons
Name
Common Formula Hybridization
Alkane
Cn H2n+2
sp3
Cycloalkane
Cn H2n
sp3
Alkene
Cn H2n
sp2
Alkyne
Cn H2n−2
sp
Aromatic
Cn H2n−6
sp2
In a hydrocarbon with n carbons, the number of hydrogens
is 2n + 2, minus 2 for each π bond or carbon ring.
Stereoisomerism
Stereoisomerism occurs at bonds such as C=C, where both
ends have two different substituents, because the rotation
of these substituents are restricted.
Cis-trans isomerism: If both ends have a hydrogen
atom substituent, then the compound exhibits cis-trans
isomerism. The cis-isomer has both hydrogen atoms on the
same side, and the trans-isomer has the hydrogen atoms
on different sides.
4
Cahn-Ingold-Prelog Rules
The CIP Rules are used to compare two substituent groups
in the E/Z and R/S groups of naming isomers.
1. Direct comparison: If the atoms that are directly connected to the stereocenter are different,
then the atom with a higher atomic number receives
higher priority.
2. Tiebreaker I: If there is a tie, then a list of atoms
two bonds away from the stereocenter is compiled
for each of the two substituent groups. The atoms
with the greatest atomic number from each list are
then compared. If they tie, then the second greatest
atoms from each list are compared. This process is
repeated until the tie is broken.
3. Tiebreaker II: If there is still a tie after consider
atoms two bonds away from the center, then atoms
three bonds away are considered in the same way in
Tiebreaker I. This process is repeated until the tie is
broken.
4. Isotopes: If two groups differ only in isotopes (and
are otherwise identical), then mass number is used
instead of atomic number in the process.
5. Double and triple bonds: If there is a double bond in the substituent group, then the double bond is treated as a bond with “ghost atoms”
(e.g. R-A=B-R’ is treated as R-(A-B)-(B-A)-R’).
Triple bonds, similarly, have two ghost atoms for
each atom.
6. Cycles: To handle a molecule containing one or
more cycles, one must first expand it into a tree
(called a hierarchical digraph by the authors) by
traversing bonds in all possible paths starting at
the stereocenter. When the traversal encounters an
atom through which the current path has already
passed, a ghost atom is generated in order to keep
the tree finite.
Criteria of Aromaticity
Functional Groups (cont.)
If a hydrocarbon
Functional Group
Name
Suffix/Prefix
R-CH2 C6 H5 (benzyl,
toluene der.
benzylBn)
R-C5 H4 N (pyridyl)
pyridine der.
pyridin-x-yl
Note: In actual compounds, change all instances of “halo”
above to halogen names (fluoro, chloro, bromo, iodo).
1. Is cyclic, i.e. possesses a carbon ring;
2. Is planar, i.e. all carbons on the ring are on the same
plane;
3. Has an uninterrupted cloud of π electrons;
4. The number of pairs of π electrons in the cloud is
an odd number, i.e. the number of π electrons in the
cloud is 4n + 2;
then the hydrocarbon is aromatic. Aromatic compounds
are highly stable (cannot undergo addition reactions), but
can undergo substitution reactions.
Functional Groups
Functional Group
R-OH (hydroxyl)
R-O-R’ (ether)
R-X (halo)
R-NH2 (amino)
R-COH (aldehyde)
R-COX (haloformyl)
R-CO-R’ (carbonyl)
R-COOH (carboxyl)
R-COO( carboxylate)
R-COO-R’ (ester)
R-CONH2 (amide)
R-CNH-R’ (ketimine)
R-CHNH (aldimine)
R-CONCO-R’
(imide)
R-N3 (azide)
R-N2 -R’ (azo)
R-OCN (cyanate)
R-NCO (isocyanate)
R-CN (nitrile)
R-NC (isonitrile)
R-NO (nitroso)
R-NO2 (nitro)
R-ONO (nitrosooxy)
R-ONO2 (nitrate)
R-SH (sulfhydryl)
R-SCN (thiocyanate)
R-NCS
(isothiocyanate)
R-CSH (carbonothioyl)
R-PH3 (phosphino)
R-C6 H5 (phenyl, Ph)
Name
alcohol
ether
haloalkane
amine
aldehyde
acyl halide
ketone
carboxylic acid
carboxylate
ester
amide
ketimine
aldimine
imide
Suffix/Prefix
-ol
ether
halo-amine
-al
-oyl halide
-one
-oic acid
-oate
-oate
-amide
iminoiminoimido-
azide
azo
cyanate
isocyanate
nitrile
isonitrile
nitroso
nitro
nitrite
nitrate
thiol
thiocyanate
isothiocyanate
azidoazocyanatoisocyanatocyanoisocyanonitrosonitronitrosooxynitroxy-thiol
thiocyanatoisothiocyanato-
thial
phosphine
benzene der.
Enantiomerism
Amino Acids
Hydrophobic amino acids:
Name
Code Name
Alanine
Ala Valine
Phenylalanine
Phe Methionine
Leucine
Leu Proline
Isoleucine
Ile Tryptophane
Code
Val
Met
Pro
Trp
Hydrophilic amino acids:
Name
Code Name
Glycine
Gly Threonine
Serine
Ser Cysteine
Tyrosine
Tyr Asparagine
Glutamine
Gln Arginine
Lysine
Lys Histidine
Aspartic acid
Asp Glutamic acid
Code
Thr
Cys
Asn
Arg
His
Glu
Protein Structure
Proteins are large biochemical complexes that contain several polypeptide compounds (amino acid chains). They
are organized into four levels of structure:
Primary structure: The chain of amino acids that make
up the protein; this chain directly controls the other levels
of protein structure.
Secondary structure: The patterns formed by segments
of the polypeptide chain; can be either α-helices or βpleated sheets.
Tertiary structure: The folding of the polypeptide to
produce a certain shape.
Quarternary structure: The geometrical bonding of
several polypeptides to form the protein.
Chirality
A molecule possessing a nonsuperimposable mirror image
is chiral.
Two mirror images of a chiral molecule are enantiomers.
A carbon that is bonded to 4 different groups is an asymmetric center. Chiral molecules have at least one asymmetric centers.
Chiral molecules rotate polarized light. Two enantiomers
-thial rotate polarized light by the same degrees, one clockwise
and one counterclockwise. A mixture of two enantiomers
-phosphane in 1:1 does not rotate polarized light, and is racemic.
phenyl-
c
2009–2011
Zee Zuo. Licensed under CC-BY-SA 3.0, United States.
System Name
R/S
(+)/(−)
Based On
Structure
Direction of rotation of polarized
light
D/L
Enantiomer of glyceraldehyde the
molecule is derived from
R/S notation: Orient the enantiomer so that the smallest (by CIP Rules) substituent points backward (away from
the viewer) and the largest substituent points upward. If
the larger substituent of the other two points toward the
right, then the enantiomer is an R-enantiomer. If the larger
substituent points toward the left, then the enantiomer is
an L-enantiomer.
(+)/(−) notation: An enantiomer that rotates the plane
of polarization clockwise is dextrorotary (+). An enantiomer that rotates the plane of polarization counterclockwise is levorotary (−).
D/L notation: An enantiomer that is derived from (+)glyceraldehyde is the D-enantiomer. An enantiomer that
is derived from (−)-glyceraldehyde is the L-enantiomer.
Note that nomenclature in a system cannot be determined
by that in another system.
Caution: The (+)/(−) system is sometimes written as
(d)/(l), which is easily confused with the D, L system. As
these two systems sometimes conflict (a D-enantiomer can
be an (l)-enantiomer), the (+)/(−) notations are strongly
preferred.
Significant Figures
Significant figures (“sig figs”) is the number of digits that
carry precision in a number.
Non-measured Numbers: Non-measured numbers,
such as π, integer counts, definition of units, etc. always
have infinite sig figs. Other constants, such as NA , have
limited sig figs.
Non-zero Digits: Nonzero digits are always significant,
unless one or more of the other rules are violated.
Zeros: Leading zeros are never significant; trailing zeros,
however, are significant only if they are part of the measurement. Zeros between non-zero digits are always significant.
Reporting Numbers: Reported numbers are only significant to the precision of the equipments with which they
are measured.
Addition/Subtraction: When adding or subtracting
two numbers, the result should have as many decimal
places as the number with the smallest sig figs.
Multiplication/Division: When multiplying or dividing, the result should have as many sig figs as the number
with the smallest sig figs.
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