• 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
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 Divalent Trivalent
Hydronium H3O+ Magnesium Mg2+ Aluminium Al3+
(or hydrogen) H+ Calcium Ca2+ Antimony III Sb3+
Lithium Li+ Strontium Sr2+ Bismuth III Bi3+
Sodium Na+ Beryllium Be2+
Potassium K+ Manganese II Mn2+
Rubidium Rb+ Barium Ba2+
Cesium Cs+ Zinc Zn2+
Francium Fr+ Cadmium Cd2+
Silver Ag+ Nickel II Ni2+
Ammonium NH4+ Palladium II Pd2+
Thalium Tl+ Platinum II Pt2+
Copper I Cu+ Copper II Cu2+
Mercury II Hg2+
Mercury I Hg22+
Iron II Fe2+ Iron III Fe3+
Cobalt II Co2+ Cobalt III Co3+
Chromium II Cr2+ Chromium III Cr3+
Lead II Pb2+
Tin II Sn2+
Table of Common Ions
Common Negative Ions (Anions)
Monovalent Divalent Trivalent
Hydride H- Oxide O2- Nitride N3-
Fluoride Fl- Peroxide O22-
Chloride Cl- Sulfide S2-
Bromide Br- Selenide Se2-
Iodide I- Oxalate C2O42-
Hydroxide OH- Chromate CrO42-
Permangante MnO4- Dichromate Cr2O72-
Cyanide CN- Tungstate WO42-
Thiocynate SCN- Molybdate MoO42-
Acetate C 2H 3O 2- tetrathionate S4O62-
Nitrate NO3- Thiosulfate S2O32-
Bisulfite HSO3- Sulfite SO32-
Bisulfate HSO4- Sulfate SO42-
Bicarbonate HCO3- Carbonate CO32-
Dihydrogen phosphate H2PO4- Hydrogen phosphate HPO42- Phosphate PO43-
or determined Calculate
from balanced stoichiometric from molecular
equation formula or balanced
mass of moles of
given or calculated from Molar Ratio
molecule periodic table molecule
Avogadro's moles of
density molarity, ppm,
molality, normality, Number element, or
etc. other reactant
Concentration Number of
given or calculated from
periodic table Number
These concepts lead to solving
problems determining limiting reactant Number of
and percent yield. atoms,
or reactant or molecules
or product of reactant
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
= 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 ↓
The Nucleus: Electrons:
Build by adding the required number of protons Hund’s Rule states that each orbital will be filled singly
(the atomic number) and neutrons (the mass of the atom) before pairing begins. The singly filled orbitals will have
a parallel spin.
Pauli’s Exclusion Principle states that paired Fill the electrons in starting with the lowest energy level
electrons in an orbital will have opposite spins. 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 Formal charge
Never Have a Full Octet Always Have a Full Octet
Sometimes Have a Full Octet
Sometimes Exceed a Full Octet
To calculate a formal charge To calculate an oxidation number
1. draw the Lewis dot structure 1. list all the elements follow with an equal sign
2. draw circles around each atom and the 2. follow with the number of atoms of that type in the
electrons associated with it. Remember that molecule
formal charges are associated with covalent 1. follow with a multiplication sign
bonds and that all electrons are shared equally. 2. If the element is O follow with a -2
3. compare to the group number for that atom. If 3. If the element is H follow with a +1
the number is larger the formal charge is 4. any other element enter a ?
negative, smaller the formal charge is positive. 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 ?
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 Electronic Geometry Hybridization
2 Linear sp
3 Trigonal planar sp2
4 Tetrahedral sp3
5 Trigonal bipyramidal sp3d
6 Octahedral sp3d2
Isomers hydrocarbons sugars fats
structural isomers unsaturated hydrocarbons polymers solution
constitutional isomers saturated hydrocarbons solvent solute
stereo isomers alkanes concentration molarity
racemic mixture alkenes ppm ppb
entantiomers alkynes wt% vol%
geometric isomers aromatic compounds molecular equations
positional isomers alkyls ionic equations
chiral molecules phenyls net ionic equations
chiral centers phenols spectator ion
optical isomers alcohols metathesis reaction
cis esters combination reaction
mer ethers decomposition reaction
trans carbonyl groups displacement reaction
fac aldehydes redox reaction
hydration isomers ketones addition polymerization
ionization isomers carboxylic acids condensation polymerization
coordination isomers acyl chlorides ligand
linkage isomers organic halides donor atom
titration amines unidentate
titrant amides polydentate
primary standard resonance chelate
secondary standard Arrhenius acids/bases coordination number
end point Brönsted/Lowery acids/bases coordination sphere
equivalence point Lewis acids/bases
oxidation numbers Non electrolytes
Naming Saturated Hydrocarbons
1. Choose the longest continuous chain of carbon atoms which gives the
basic name or stem.
2 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
3 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.
4 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
5 The combined substituent numbers and names serve as a prefix for the
basic hydrocarbon name.
6 Separate numbers from numbers by commas and numbers from words
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.
Aldehydes and Ketones
• Common names for aldehydes are derived from the name of the acid with the same number of C
• 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
• A numeric prefix indicates the position of the carbonyl group in a chain or on a ring.
• 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.
• 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
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
Cations -onio- -onium
e.g. Ammonium –NH4+ ammonio- -ammonium
2 Carboxylic acids –COOH carboxy- -oic acid*
Carboxylic acid derivatives
Esters –COOR R-oxycarbonyl-
Acyl chlorides –COCl chloroformyl- -oyl chloride*
Amides –CONH2 carbamoyl- -amide*
Nitrites –CN cyano- -nitrile*
Isocyanides –NC isocyano- isocyanide
Aldehydes –CHO formyl- -al*
Thioaldehydes –CHS thioformyl- -thial*
Ketones >CO oxo- -one
Thioketones >CS thiono- -thione
Alcohols –OH hydroxy- -ol
Thiols –SH sulfanyl- -thiol
8 Amines –NH2 amino- -amine
Ethers –O– -oxy-
Thioethers –S– -thio-
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
Rules for Naming Complex Species
1. Cations (+ ions) are named before anions (- ions).
2. 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.
3. 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
4. The names of most anionic ligands end in the suffix -o.
– Examples of ligands ending in –o are:
• Cl- chloro S2- sulfido O2- oxo
5. 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
6. 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.
7. 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 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)
heat lost or gained by system (Joules)
Specific heat capacity (J/(g∙K) =
mass(grams) DT (Kelvins)
Variabl System System
q e 1 2
m(Tf –Ti) Tf
heat transfer in heat transfer out
(endothermic), +q (exothermic), -q
∆E = q + w
w transfer in w transfer out
State Function Vapor Pressure
Standard state temperature Equilibrium
Standard state pressure Heat of Vaporization
Standard states matter Phase Diagram
Hess’s Law Liquid
Thermochemical Equation Gas
Enthalpy of Formation Triple Point
Intramolecular forces Critical Point
Intermolecular forces Super Critical Fluid
Law of Conservation of Energy
Specific Heat Capacity
First Law of Thermodynamics
Standard P 1.00000 atm or 101.3 kPa n 2a
Standard T 273.15 K or 0.00oC P + V nb nRT
K = 273 + oC
1 mm Hg = 1 torr 760 torr = 1 atm Variable Cond. Cond.
The standard molar volume is 22.4 L at STP P (atm)
PV = nRT
R = 0.08206 L atm mol-1 K-1 R (L atm mol-1 K-1) 0.08206 0.08206
Ptotal = PA + PB + PC + .....
At low temperatures and high pressures real gases do not behave
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.
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
• Proof - Gases are easily compressible.
• Postulate 2
– Gas molecules are in constant, random, straight line motion with varying
• 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
– The average kinetic energies of molecules of different gases are equal at a
• Proof - Brownian motion increases as temperature increases.