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•
Significant Figures – start at the left and proceed to the right
1. If the number does not have a decimal point count until there are no
more non zero numbers
2. If the number has a decimal point start counting at the first non-zero
number and continue counting until you run out of decimal places
• Vocabulary
1. Observation
17. Law of Conservation of Mass 33. percent weight
2. Hypothesis
18. Law of Conservation of Energy 34. percent error
3. Experiment
19. Exact numbers
35. percent composition
4. Theory
20. Accuracy
36. percent yield
5. Law
21. Precision
37. %RSD
6. Chemistry
22. compounds
38. limiting reactant
7. Matter
23. molecules
39. Stoichiometry
8. Energy
24. chemical formula
40. Stoichiometric Coefficient
9. Chemical Properties 25. empirical formula
41. Electron Affinity
10. Physical Properties
26. molecular formula
42. Electronegativity
11. Extensive Properties 27. structural formula
43. Covalent Bond
12. Intensive Properties
28. bond line formula
44. Ionic Bond
13. Scientific (natural) law 29. ball and stick model
45. Dipole
14. Anion
30. space filling model
46. London Dispersion Forces
15. Cation
31. mole
47. Resonance
16. Molecular Geometry 32. Electronic Geometry
48. Hybrid orbital
49. area of high electron density
Table of Common Ions
Common Positive Ions (Cations)
Monovalent
Hydronium
(or hydrogen)
Lithium
Sodium
Potassium
Rubidium
Cesium
Francium
Silver
Ammonium
Thalium
Copper I
H3O+
H+
Li+
Na+
K+
Rb+
Cs+
Fr+
Ag+
NH4+
Tl+
Cu+
Divalent
Magnesium
Calcium
Strontium
Beryllium
Manganese II
Barium
Zinc
Cadmium
Nickel II
Palladium II
Platinum II
Copper II
Mercury II
Mercury I
Iron II
Cobalt II
Chromium II
Lead II
Tin II
Mg2+
Ca2+
Sr2+
Be2+
Mn2+
Ba2+
Zn2+
Cd2+
Ni2+
Pd2+
Pt2+
Cu2+
Hg2+
Hg22+
Fe2+
Co2+
Cr2+
Pb2+
Sn2+
Trivalent
Aluminium
Antimony III
Bismuth III
Al3+
Sb3+
Bi3+
Iron III
Cobalt III
Chromium III
Fe3+
Co3+
Cr3+
Table of Common Ions
Common Negative Ions (Anions)
Monovalent
Hydride
Fluoride
Chloride
Bromide
Iodide
Hydroxide
Permangante
Cyanide
Thiocynate
Acetate
Nitrate
Bisulfite
Bisulfate
Bicarbonate
Dihydrogen phosphate
Nitrite
Amide
Hypochlorite
Chlorite
Chlorate
Perchlorate
HFlClBrIOHMnO4CNSCNC 2H 3O 2NO3HSO3HSO4HCO3H2PO4NO2NH2ClOClO2ClO3ClO4-
Divalent
Oxide
Peroxide
Sulfide
Selenide
Oxalate
Chromate
Dichromate
Tungstate
Molybdate
tetrathionate
Thiosulfate
Sulfite
Sulfate
Carbonate
Hydrogen phosphate
Trivalent
Nitride
O2O22S2Se2C2O42CrO42Cr2O72WO42MoO42S4O62S2O32SO32SO42CO32HPO42- Phosphate
N3-
PO43-
Given
or determined
from balanced stoichiometric
equation
mass of
molecule
Molar Mass
given or calculated from
periodic table
density
molarity, ppm,
molality, normality,
etc.
Vol solution
Concentration
solution
Calculate
from molecular
formula or balanced
equation
moles of
molecule
Avogadro's
Number
Molar Ratio
moles of
element, or
other reactant
or product
Number of
molecules
These concepts lead to solving
problems determining limiting reactant
and percent yield.
Molar Mass
given or calculated from
periodic table
Avogadro's
Number
Mass of
element,
or reactant
or product
Number of
atoms,
or molecules
of reactant
or product
Quantum Numbers
n and l define the energy of the electron
The principal quantum number has the symbol ~ n which defines the
energy of the shell
n = 1, 2, 3, 4, ...... “shells”
The angular momentum quantum number has the symbol ~  which defines the
subshells.
 = 0, 1, 2, 3, 4, 5, .......(n-1)
 = s, p, d, f, g, h, .......(n-1)
The symbol for the magnetic quantum number is m which defines the orbital.
m = -  , (-  + 1), (-  +2), .....0, ......., ( -2), ( -1), 
The last quantum number is the spin quantum number which has the symbol m s which characterizes the single electron.
The spin quantum number only has two possible values. ms = +½ or -½ one spin up ↑ and one spin down ↓
Electrons:
The Nucleus:
Hund’s Rule states that each orbital will be filled singly
Build by adding the required number of protons
before pairing begins. The singly filled orbitals will have
(the atomic number) and neutrons (the mass of the atom)
a parallel spin.
Pauli’s Exclusion Principle states that paired
electrons in an orbital will have opposite spins.
Fill the electrons in starting with the lowest energy level
adhering to Hund’s and Pauli’s rules.
Ionic
Polar Covalent
Covalent
Determine Inductive effect
Count the number of electrons the element should have
Determine how equally electrons are shared (DEN) >1.7 consider it ionic
Oxidation number
Never Have a Full Octet
Formal charge
Always Have a Full Octet
Sometimes Have a Full Octet
Sometimes Exceed a Full Octet
To calculate a formal charge
1. draw the Lewis dot structure
2. draw circles around each atom and the
electrons associated with it. Remember that
formal charges are associated with covalent
bonds and that all electrons are shared equally.
3. compare to the group number for that atom. If
the number is larger the formal charge is
negative, smaller the formal charge is positive.
To calculate an oxidation number
1. list all the elements follow with an equal sign
2. follow with the number of atoms of that type in the
molecule
1. follow with a multiplication sign
2. If the element is O follow with a -2
3. If the element is H follow with a +1
4. any other element enter a ?
5. follow with an = sign, do the math
6. draw a total line, then enter the charge on the molecule
7. Do the algebra backwards to solve for ?
VSEPR Theory
Lone pair to lone pair is the strongest repulsion.
2 Lone pair to bonding pair is intermediate repulsion.
3 Bonding pair to bonding pair is weakest repulsion.
•
Mnemonic for repulsion strengths
lp/lp > lp/bp > bp/bp
•
Lone pair to lone pair repulsion is why bond angles in water are less than 109.5o.
Electronic geometry is determined by the locations of regions of high electron density around the central
atom(s). Electron pairs are not used in the molecular geometry determination just the positions of
the atoms in the molecule are used.
Molecular geometry determined by the arrangement of atoms around the central atom(s).
Summary of Electronic & Molecular Geometries
Regions of High Electron
Density
Electronic Geometry
Hybridization
2
Linear
sp
3
Trigonal planar
sp2
4
Tetrahedral
sp3
5
Trigonal bipyramidal
sp3d
6
Octahedral
sp3d2
Isomers
structural isomers
constitutional isomers
stereo isomers
racemic mixture
entantiomers
geometric isomers
positional isomers
chiral molecules
chiral centers
optical isomers
cis
mer
trans
fac
hydration isomers
ionization isomers
coordination isomers
linkage isomers
titration
titrant
primary standard
secondary standard
end point
equivalence point
pH
oxidation numbers
hydrocarbons
unsaturated hydrocarbons
saturated hydrocarbons
alkanes
alkenes
alkynes
aromatic compounds
alkyls
phenyls
phenols
alcohols
esters
ethers
carbonyl groups
aldehydes
ketones
carboxylic acids
acyl chlorides
organic halides
amines
amides
resonance
Arrhenius acids/bases
Brönsted/Lowery acids/bases
Lewis acids/bases
Electrolytes
Non electrolytes
sugars
polymers
solvent
concentration
ppm
wt%
molecular equations
ionic equations
net ionic equations
spectator ion
metathesis reaction
combination reaction
decomposition reaction
displacement reaction
redox reaction
addition polymerization
condensation polymerization
ligand
donor atom
unidentate
polydentate
chelate
coordination number
coordination sphere
fats
solution
solute
molarity
ppb
vol%
Naming Saturated Hydrocarbons
1.
2
3
4
5
6
Choose the longest continuous chain of carbon atoms which gives the
basic name or stem.
Number each carbon atom in the basic chain, starting at the end that
gives the lowest number to the first group attached to the main chain
(substituent).
For each substituent on the chain, we indicate the position in the chain
(by an Arabic numeric prefix) and the kind of substituent (by its name).
 The position of a substituent on the chain is indicated by the lowest
number possible. The number precedes the name of the substituent.
When there are two or more substituents of a given kind, use prefixes to
indicate the number of substituents.
 di = 2, tri = 3, tetra = 4, penta = 5, hexa = 6, hepta = 7, octa = 8, and
so on.
The combined substituent numbers and names serve as a prefix for the
basic hydrocarbon name.
Separate numbers from numbers by commas and numbers from words
by hyphens.
 Words are "run together".
Alcohols and Phenols
•
•
•
The stem for the parent hydrocarbon plus an -ol suffix is the systematic name for an alcohol.
A numeric prefix indicates the position of the -OH group in alcohols with three or more C atoms.
Common names are the name of the appropriate alkyl group plus alcohol.
•
Common names are used for most ethers.
Ethers
Aldehydes and Ketones
•
•
•
•
Common names for aldehydes are derived from the name of the acid with the same number of C
atoms.
IUPAC names are derived from the parent hydrocarbon name by replacing -e with -al.
The IUPAC name for a ketone is the characteristic stem for the parent hydrocarbon plus the suffix
-one.
A numeric prefix indicates the position of the carbonyl group in a chain or on a ring.
Amines
•
•
Amines are derivatives of ammonia in which one or more H atoms have been replaced by
organic groups (aliphatic or aromatic or a mixture of both).
There are three classes of amines.
Carboxylic Acids
•
•
IUPAC names for a carboxylic acid are derived from the name of the parent hydrocarbon.
– The final -e is dropped from the name of the parent hydrocarbon
– The suffix -oic is added followed by the word acid.
Many organic acids are called by their common (trivial) names which are derived from Greek or
Latin.
When compounds contain more than one functional group, the order of precedence determines
which groups are named with prefix or suffix forms. The highest precedence group takes the
suffix, with all others taking the prefix form. However, double and triple bonds only take suffix
form (-en and -yn) and are used with other suffixes.
Functional group
Formula
Prefix
Suffix
1
Cations
e.g. Ammonium
–NH4+
-onioammonio-
-onium
-ammonium
2
Carboxylic acids
–COOH
carboxy-
-oic acid*
3
Carboxylic acid derivatives
Esters
Acyl chlorides
Amides
–COOR
–COCl
–CONH2
R-oxycarbonylchloroformylcarbamoyl-
-oyl chloride*
-amide*
4
Nitrites
Isocyanides
–CN
–NC
cyanoisocyano-
-nitrile*
isocyanide
5
Aldehydes
Thioaldehydes
–CHO
–CHS
formylthioformyl-
-al*
-thial*
6
Ketones
Thioketones
>CO
>CS
oxothiono-
-one
-thione
7
Alcohols
Thiols
–OH
–SH
hydroxysulfanyl-
-ol
-thiol
8
Amines
–NH2
amino-
-amine
9
Ethers
Thioethers
–O–
–S–
-oxy-thio-
Priority
Carbon Atom Hybridization
C uses
C forms
Example
sp3 tetrahedral
4 sp3 hybrids
4  bonds
CH4
sp2 trigonal planar
3 sp2 hybrids & 1p orbital
3  bonds 1  bond
C2H4
sp linear
2 sp hybrids & 2 p orbitals
2  bonds 2  bonds
C2H2
Nomenclature
1.
2.
3.
4.
5.
6.
7.
Rules for Naming Complex Species
Cations (+ ions) are named before anions (- ions).
Coordinated ligands are named in alphabetical order.
–
Prefixes that specify the number of each kind of ligand (di = 2, tri = 3, tetra = 4, penta =
5, hexa = 6, etc.) are not used in alphabetizing
–
Prefixes that are part of the name of the ligand, such as in diethylamine, are used to
alphabetize the ligands.
For complicated ligands, especially those that have a prefix such as di or tri as part of the
ligand name, these prefixes are used to specify the number of those ligands that are
attached to the central atom.
–
bis = 2 tris = 3 tetrakis = 4 pentakis = 5 hexakis = 6
The names of most anionic ligands end in the suffix -o.
–
Examples of ligands ending in –o are:
•
Cl- chloro
S2- sulfido
O2oxo
The names of most neutral ligands are unchanged when used in naming the complex.
–
There are several important exceptions to this rule including:
•
NH3 ammine
H2O aqua
The oxidation number of a metal that exhibits variable oxidation states is designated by a
Roman numeral in parentheses following the name of the complex ion or molecule.
If a complex is an anion, the suffix "ate" ends the name.
 No suffix is used in the case of a neutral or cationic complex.
 Usually, the English stem is used for a metal, but if this would make the name awkward,
the Latin stem is substituted. ferrate instead of ironate plumbate instead of leadate
Ion/Molecule
Name
Name as a Ligand
NH3
ammonia
ammine
CO
carbon monoxide
carbonyl
Cl-
chloride
Chloro
CN-
cyanide
cyano
F-
fluoride
fluoro
OH-
hydroxide
hydroxo
NO
nitrogen monoxide
nitrosyl
NO2-
nitrite
nitro
PH3
phosphine
phosphine
System
Arrhenius
BrönstedLowry
Lewis
Acid (HCl)
Base (NaOH)
∆H = Hfinal - Hinitial
C5 H12( )  8 O 2(g)  5 CO2(g)  6 H 2 O (  )  3523 kJ
1 mole
•
8 moles
5 moles
6 moles
1 mole
The stoichiometric coefficients in thermochemical equations must be interpreted as
numbers of moles. 1 mol of C5H12 reacts with 8 mol of O2 to produce 5 mol of CO2, 6
mol of H2O, and releasing 3523 kJ is referred to as one mole of reactions.
∆Horxn =  ∆Hfo (prod) -  ∆Hfo (react)
Specific heat capacity (J/(g∙K) =
cP =
q
m(Tf –Ti)
heat lost or gained by system (Joules)
mass(grams) DT (Kelvins)
Variabl
e
Cp
Tf
Ti
m
q
System
1
System
2
heat transfer in
(endothermic), +q
heat transfer out
(exothermic), -q
SYSTEM
∆E = q + w
w transfer in
(+w)
w transfer out
(-w)
Heat
Energy
Internal energy
Kinetic Energy
Potential Energy
Endothermic
Exothermic
Thermodynamics
Thermal Equilibrium
System
Surroundings
Law of Conservation of Energy
Heat Capacity
Specific Heat Capacity
First Law of Thermodynamics
Melting
Freezing
Deposition
Sublimation
Evaporation
Condensation
State Function
Standard state temperature
Standard state pressure
Standard states matter
Enthalpy
Hess’s Law
Thermochemical Equation
Enthalpy of Formation
Intramolecular forces
Intermolecular forces
Hydrogen Bonding
Polarization
Polarizability
Vapor Pressure
Equilibrium
Heat of Vaporization
Phase Diagram
Solid
Liquid
Gas
Triple Point
Critical Point
Super Critical Fluid
Standard P  1.00000 atm or 101.3 kPa
Standard T  273.15 K or 0.00oC

n 2a 
V  nb  nRT
P +
2 (
V 

K = 273 + oC
1 mm Hg = 1 torr 760 torr = 1 atm
Variable
The standard molar volume is 22.4 L at STP
P (atm)
PV = nRT
R = 0.08206 L atm mol-1 K-1
Cond.
1
Cond.
2
0.08206
0.08206
V (L)
N (moles)
R (L atm mol-1 K-1)
T (K)
Ptotal = PA + PB + PC + .....
At low temperatures and high pressures real gases do not behave
ideally.
The reasons for the deviations from ideality are:
1. The molecules are very close to one another, thus their
volume is important.
2. The molecular interactions also become important.
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•
The Kinetic-Molecular Theory
The basic assumptions of kinetic-molecular theory are:
Postulate 1
– Gases consist of discrete molecules that are relatively far apart.
– Gases have few intermolecular attractions.
– The volume of individual molecules is very small compared to the gas’s
volume.
Proof - Gases are easily compressible.
Postulate 2
– Gas molecules are in constant, random, straight line motion with varying
velocities.
Proof - Brownian motion displays molecular motion.
Postulate 3
– Gas molecules have elastic collisions with themselves and the container.
– Total energy is conserved during a collision.
Proof - A sealed, confined gas exhibits no pressure drop over time.
Postulate 4
– The kinetic energy of the molecules is proportional to the absolute
temperature.
– The average kinetic energies of molecules of different gases are equal at a
given temperature.
Proof - Brownian motion increases as temperature increases.