JOSH’S
FANTASTERIFFIC
CHEMISTRY
REVIEW PACKET!
VOCABULARY:
You should recognize/know all of the following terms, be able to describe them, use them in a sentence, and
understand their connection to other chemistry terms and chemistry as a whole.
This is a list of every vocab word in the book (yes, you should know them all). Go through the list and
HIGHLIGHT every word you do not recognize, cannot define, or do not know the meaning of in terms of
chemistry. Then, using the space to the right, define the words.
Chapter 2






Density
Dimensional Analysis
Accuracy
Error
Percent Error
Significant Figure






















Chemical Property
Extensive Property
Intensive Property
Gas
Liquid
Solid
Physical Property
States of Matter
Vapor
Chemical Change
Law of conservation of mass
Physical change
Phase Change
Heterogeneous Mixture
Homogeneous Mixture
Mixture
Solution
Compound
Element
Law of Definite Proportions
Law of multiple proportions
Percent mass










Dalton’s Atomic theory
Atom
Electron
Neutron
Nucleus
Proton
Atomic Mass
Isotope
Alpha Particle
Alpha Radiation
Beta particle/radiation
Gamma Ray
Nuclear Equation
Radioactive decay



















Amplitude
Electromagnetic Radiation
Frequency
Photon
Planck’s Constant
Quantum
Wavelength
Atomic Orbital
De Broglie equation
Energy sublevel
Ground State
Heisenberg Uncertainty Principle
Quantum Number
Aufbau Principle
Electron Configuration
Electron-dot Structure
Hund’s Rule
Pauli Exclusion Principle
Valence Electron















Actinide Series
Alkali metal
Alkaline earth metal
Group
Halogen
Inner Transition Metal
Lanthanide Series
Metal
Metalloid
Noble Gas
Nonmetal
Period
Periodic Law
Representative Element
Transition Element
Chapter 5
Chapter 3
Chapter 4




Chapter 6





Transition Metal
Electronegativity
Ion
Ionization energy
Octet Rule

















Anion
Cation
Chemical Bond
Crystal Lattice
Electrolyte
Ionic Bond
Ionic Compound
Lattice Energy
Formula Unit
Monatomic Ion
Oxidation Number
Oxyanion
Polyatomic Ion
Alloy
Delocalized Electron
Electron Sea Model
Metallic Bond










Covalent Bond
Endothermic Reaction
Exothermic Reaction
Lewis Structure
Molecule
Pi Bond
Sigma Bond
Resonance
Structural formula
Polar Covalent Bond














Chemical Equation
Chemical Reaction
Coefficient
Product
Reactant
Combustion Reaction
Decomposition Reaction
Double-Replacement Reaction
Precipitate
Single-Replacement Reaction
Synthesis Reaction
Aqueous Solution
Complete Ionic Equation
Net Ionic equation
Chapter 10
Chapter 7
Chapter 8
Chapter 9



Solute
Solvent
Spectator Ion







Avogadro’s Number
Mole
Molar Mass
Empirical Formula
Molecular Formula
Percent Composition
Hydrate








Mole Ratio
Stoichiometry
Excess Reactant
Excess Reactant
Limiting Reactant
Actual yield
Percent Yield
Theoretical Yield



























Atmosphere
Barometer
Dalton’s Law of Partial Pressures
Diffusion
Elastic Collision
Graham’s Law of effusion
Kinetic-Molecular Theory
Pascal
Pressure
Dipole-Dipole Force
Hydrogen Bond
Dispersion Force
Amorphous Solid
Crystalline Solid
Surfactant
Surface Tension
Unit Cell
Viscosity
Boiling Point
Condensation
Deposition
Evaporation
Freezing Point
Melting Point
Phase Diagram
Triple Point
Vaporization
Chapter 11
Chapter 12

Vapor Pressure









Absolute Zero
Boyle’s Law
Charles’s Law
Combined Gas Law
Gay-Lussac’s Law
Avogadro’s Principle
Ideal gas Constant
Ideal Gas Law
Molar Volume
Chapter 13
Chapter 14









Solvation
Supersaturated Solution
Unsaturated Solution
Boiling Point Elevation
Colligative Property
Freezing Point depression
Osmosis
Osmotic Pressure
Vapor Pressure Lowering


















Calorie
Chemical Potential Energy
Energy
Heat
Joule
Law of Conservation of Energy
Calorimeter
Enthalpy
Enthalpy of reaction
Surroundings
System
Universe
Thermochemistry
Enthalpy of Combustion
Molar Enthalpy of Fusion
Molar Enthalpy of Vaporization
Thermochemical Equation
Hess’s Law
Chapter 15
 Brownian Motion
 Colloid
 Immisible
 Insoluble
 Miscible
 Soluble
 Suspension
 Tyndall Effect
 Concentration
 Molality
 Molarity
 Mole Fraction
 Heat of Solution
 Henry’s Law
 Saturated Solution
**Vocabulary from the final chapters is not included..mostly because I don’t know what chapters we’re going
to do! HOPEFULLY you won’t have to review the ending chapters very much since they are so close to the
final exam, but the vocabulary is listed at the end of the chapters if you need to. 
Chemistry Joke: Why do chemists like nitrates so much? They are cheaper than day rates. LOL!!!
CHAPTER CONCEPT SUMMARIES
Below are chapter summaries, which include the main ideas and concepts included in each chapter. Scan them
over and make sure you understand the material, but keep in mind this is the basic information. There is no
substitute for the book! If you come across some material you do not understand, or do not remember, or even
have the slightest doubt about, go back to the book and re-READ that section! (don’t just look up the
information, read the entire section. It helps!)
CHAPTER 2: Analyzing Data
 Density is mass / volume. [grams/milliliter]
 Be able to recognize general conversion factors (like hours in a minute or days in a year, or number of
molecules in a mole…)
 Be able to use dimensional analysis to change units.
o Note from Josh: It is important to recognize the units on different numbers. If you are working
the problem and have no idea what to do, but can keep the units straight, you will usually be able
to figure things out!

Accuracy refers to how close a measured value is to an accepted value. Precision refers to how close a
series of measurements are to one another.
 Percent Error = [(experimental value – accepted value)/accepted value] * 100
 SIGNIFIGANT FIGURE RULES
o Nonzero numbers are always significant
o Zeros between nonzero numbers are always significant
o All final zeros to the right of the decimal are significant
o Placeholder Zeroes are not significant. To remove placeholder zeroes, rewrite the number in
scientific notation.
o Counting numbers are defined as constants and have an infinite number of significant figures.
Chapter 3: Matter – Properties and Changes
 Solids have their own definite shape and volume. Liquids have a constant volume and take the shape of
their container. Gases also have constant volume and take the shape of their container, but they will flow
to fill the container as a whole.
 Physical Property : a characteristic that can be observed without changing the sample’s composition.
(i.e. density, color, boiling/melting point)
 Extensive Properties depend on the amount of the substance present (mass for example). Intensive
properties do not (like density).
 Chemical Property: the ability of a substance to combine with or change into one or more other
substances.
 LAW OF CONSERVATION OF MASS: mass is neither created nor destroyed during a chemical
reaction.
 Mixture
o Homogeneous: constant composition throughout
o Heterogeneous: different throughout
 Solution: another name for a homogeneous mixture.
 LAWOF DEFINITE PROPORTIONS: a compound is always composed of the same elements in the
same proportion by mass, no matter how large or small the sample. Percent mass = (mass of element /
mass of compound) * 100
 LAW OF MULTIPLE PROPORTIONS: when different compounds are formed by a combination of the
same elements, different masses of one element combine with the same relative mass of the other
element in a ratio of small whole numbers.
Chapter 4: The Structure of the Atom
 Dalton’s Atomic Theory:
o Matter is composed of extremely small particles called atoms.
o Atoms are indivisible and indestructible
o Atoms of a given element are identical in size, mass, and chemical properties.
o Atoms of a specific element are different from those of another element.
o Different atoms combine in simple, whole-number ratios to form compounds.
o In a chemical reaction, atoms are separated, combined or rearranged.
 Note from Josh: You should be able to recognize the general history behind different models of the atom.
I advise to go back and scan this section, 106-115, and at least be able to recognize different models and
experiments. I doubt any extreme historical knowledge will be necessary. :D
 Isotopes of elements exist: they have the same number of protons but different numbers of electrons.
 Check out table 4.5 on page 124! It has the radiation particles. (alpha, beta, and gamma radation)
Chapter 5: Electrons in Atoms
 The Wave nature of light:
o Wavelength (lambda, the upside down y shape) is the shortest distance between two equivalent
points on a wave
o Frequency (nu, which looks like a v) is the number of waves that pass a given point per second.
o Amplitude is the length from the waves height from the origin to a crest or from the origin to a
trough.
o Electromagnetic wave relationship C = (wavelength)(frequency), where C is the speed of light,
3E8
 A quantum is the minimum amount of energy that can be gained or lost by an atom.
o Equantum = hv, where v is the frequency and h is plank’s constant, 6.626E-34
 The lowest possible energy level an atom can have is called its ground state. Atoms can gain energy,
measured in quanta, and begin to enter an exited state.
 Electron configuration!!!
o The S orbitals should be paired with the first two families on the periodic table. (S=2)
o The P orbitals should be paired with the last 6. (P=6)
o The D orbitals should be paired with the transition metals. (d=10)
o The F orbitals should be paired with the lanthanide and actinides. (at the bottom of the periodic
table) (f = 14)
o To write out the electron configuration for a particular element, start at the top left of the periodic
table, with hydrogen, and then go through each element, noting its configuration mentally, until
you get to the element you want. You can also reference a noble gas, by placing that noble gas in
brackets and then moving on from there.
 Electron Configuration Rules:
o Afbau Principle: each electron occupies the lowest energy orbital available.
o Pauli Exclusion Principle: a maximum of two electrons can occupy a single atomic orbital, but
only if they have opposite spins.
o Hund’s Rule: single electrons with the same spin must occupy each equal-energy orbital before
additional electrons with opposite spins can occupy the same orbitals.
 The electron dot structure includes the elemental symbol
Chapter 6: The Periodic Table and Periodic Law:
 Recognize the general history of the periodic table regarding its discovery (*Mendeleev) but no extreme
historical knowledge will be necessary.
 Periods = Rows; Groups = columns.
o Group 1- Alkali Metals
o Group 2- Alkaline Earth Metals
o Transition Metals: middle of the periodic table
o Group 17- Halogens
o Group 18- Noble gases
 Periodic Trends:
o Atomic Radius: Large on bottom left of periodic table, small on top right of periodic table.
o Ionic Radius: note that this is the radius of an ION, not an atom! Generally these increase when
going down a group, but look at the picture on page 190 for a better view.
o Ionization energy: the energy required to remove an electron from a gaseous atom. Ionization
energy is generally larger on the top right of the periodic table and smaller on the bottom left of
the periodic table.
o Electronegativity: the relative ability of its atoms to attract electrons in a chemical bond.
Electronegativity is generally smaller on the bottom left of the periodic table and greater on the
top right of the periodic table.
 **OCTET RULE: atoms tend to gain, lose, or share electrons in order to acquire a full set of eight
valence electrons.
Chapter 7: Ionic Compounds and Metals
 A positively charged ion is a cation. A negatively charged ion is an anion.
 Ionic Compounds consist of a metal and a nonmetal. Remember that in an ionic compound, a crystalline
structure is formed between the ions, and a crystal lattice, a geometric arrangement of particles, is
formed. Melting point, boiling point, and hardness depend on how strongly the ions are attracted to one
another
 Lattice energy: the strength of the forces holding ions in place.
 Naming Ionic Compounds:
o The metal comes first, followed by the nonmetal and the appropriate ending. This will be ide
most of the time, unless a polyatomic ion is being used. Remember that sometimes transition
metals will be used for the metal, and since they can have more than one charge, the charge
needs to be specific. (example: FeO is Iron (II) Oxide, but Fe2O3 is Iron (III) Oxide.
 Electron Sea Model: all the metal atoms in a metallic solid contribute to their valence electrons to form a
‘sea’ of electrons. Because they are free to move, they are often referred to a delocalized electrons.
 Metals
o Melting/boiling points: great variation
o Malleability, ductility, durability. (can be hammered into sheets, drawn into wire, and they last a
long time).
o Thermal conductivity and electrical conductivity.
o Hardness and strength
o An alloy is a combination of elements that have metallic properties. (brass, stainless steel, and
cast iron are various alloys).
 In substitutional alloys, some of the atoms in the original metallic solid are replaced by
other metals of similar atomic size.
 In interstitial alloys, small holes of the metallic crystal are filled with smaller atoms.
Chapter 8: Covalent Bonding:
 A covalent bond consists of a nonmetal and a nonmetal. Electrons are shared rather than transferred
 A sigma bond occurs when the pair of shared electrons is in an area centered between the two atoms. (in
a lewis dot structure, a single line is a sigma bond)
 A pi bond forms when parallel orbitals overlap and share electrons. A double bond is one sigma bond
and one pi bond; a triple bond is one sigma bond and 2 pi bonds.
Note from Josh: When naming compounds, identify what kind of compound it is first, and then go from
there. Use what you know!
 Resonance is a condition that occurs when more than one valid Lewis structure can be written for a
molecule or ion.
 A polar covalent bond is formed when there is an unequal sharing of electrons.
Chapter 9: Chemical Reactions:
 The process by which atoms of one or more substances are rearranged to form a different substance is
called a chemical reaction. (Be sure to balance equations! Whenever you are doing a problem involving
a chemical equation it is probably wise to make sure that the equation is balanced.)
 Types of Reactions
o Synthesis: A + B  AB
o Combustion: oxygen combines with a substance and releases energy in the form of heat and light.
Whenever an organic material combusts, the products are ALWAYS carbon dioxide and water.
o Decomposition: AB  A + B
o Single Replacement: A + BX  AX + B
o Double Replacement: AX + BY  AY + BX
 **know how to write out ionic equations. (complete equation, net equation. Etc.)
Chapter 10: The Mole:
 If you don’t know what a mole is by now then you are in some trouble
 Percent composition = (mass of element) / (mass of compound) times 100
 LOOK AT EXAMPLE 10.10 on page 343. You should know how to do a problem like that
Chemistry Joke: Q: If H-two-O is the formula for water, what is the formula for ice?
A: H-two-O-CUBED. LOL!!!! ;-)
Chapter 11: Stoichiometry
 Stoichiometry is the study of quantitative relationships between the amounts of reactants used and the
amounts of products formed. It is based on the law of conservation of mass. REMEMBER TO HAVE A
BALANCED CHEMICAL EQUATION!!! The rest is basically dimensional analysis
 Limiting Reactants
o The limiting reactant limits the reaction and the excess reactant consists of reactants left over
after the reaction has stopped.
o Check out the sample problems to see how these work.
o Theoretical Yield is what should happen according to calculation, and actual yield is the actual
number received in an experiment or lab,
Chapter 12: States of Matter:
 Gases
o Have general knowledge of the kinetic molecular theory (pg. 403)
o Diffusion: the movement of one material through another. Effusion: a gas escapes through a tiny
opening.
o Grahams law of effusion/diffusion: RateA/RateB = square root (molar mass B / molar mass A)
o Pressure is defined as force per unit area.
 Pascal
 Kilopascal
 Atmosphere
 mmHG
 Torr
 1 atm = 760 mmHg = 101.3 kPa = 760 torr
o Dalton’s law of Partial Pressures: the total pressure of a mixture of gases is equal to the sum of
the pressures of all the gases in the mixture.
 FORCES OF ATTRACTION
o Dispersion Forces: weak forces that result from the temporary shifts in the density of electrons in
the electron clouds. Dispersion forces are sometimes called London forces.
o Dipole-Dipole forces: attractions between oppositely charged regions of polar molecules.
o Hydrogen Bond: a special type of dipole-dipole bond; attraction that occurs between molecules
bonded to a small, highly electronegative atom with at least one electron pair. Hydrogen bonds
are typically the strongest.
 Liquids
o Take the shape of their container
o Viscosity: a measure of the resistance of a liquid to flow. Typically, the stronger the
intermolecular forces, the higher the viscosity.
o Surface tension: the energy required to increase the surface area of a liquid by a given amount. A
measure of the inward pull by particles in the interior.
o Compounds that lower the surface tension of water all called surfaceactive agents, or surfactants.
 Solids
o Have crystalline shapes. Recognize table 12.5 on page 423
 Recognize and be able to detail phase diagrams! (page 429)
Chapter 13: Gases
 Boyles Law: P1V1 = P2V2
 Charles Law: V1/T1 = V2/T2
 Gay-Lussacs’ law: P1/T1 = P2/T2
 Ideal Gas Law: PV = nRT
Note from Josh: The ideal Gas law can be used in most any case involving gases, but it is important to know
the other laws in the event that you aren’t given enough information to use the ideal gas law. If you
somehow forget one of the gas law equations, just think of the bumper car metaphor!
 The ideal gas constant is .0821 l*atm/mol*k
 The ideal gas equation can be rearranged to find molar mass and also to find density. (pg 458)
Chapter 14: Mixtures and Solutions
 Suspension: a mixture containing particles that settle out if left undisturbed.
 Colloid: particles do not settle out and the particles are between 1nm and 1000 nm. (ex. Milk) The
dispersed particles of liquid colloids make jerky, random movements, called Brownian Motion.
 Tyndall effect: the scattering of light, due to particles that exist in a solution such as a colloid.
 Solutions:
o Soluble: a substance that is able to dissolve in a solvent is said to be soluble. If it cannot, then it
is insoluble.
o Miscible: liquids that are soluble in each other in any proportion. Two liquids that can be mixed
together but separate shortly after are said to be immiscible.
 EXPRESSING CONCENTRATION:
o Percent by mass: mass of solute over mass of solution times 100
o Percent by volume: volume of solute over volume of solution times 100
o Molarity: moles of solute over liters of solution
o Molality: moles of solute over kilograms of solvent
o Mole fraction: moles of solute over moles of solute + moles
 Solvation: the process of surrounding solute particles with solvent particles to form a solution.
o Factors that affect salvation include agitation, surface area, and temperature.
 Heat of Solution: the overall energy change that occurs during the solution formation process.
 Solubility
o Unsaturated Solution: one that contains less dissolved solute for a given temperature and
pressure than a saturated solution.
o Saturated Solution: contains the maximum amount of dissolved solute for a given amount of
solvent at a specific temperature and pressure.
o Supersaturated Solution: contains more dissolved solute than a saturated solution at the same
temperature.
 Colligative Properties: physical properties of solutions that are affected by the number of particles but
not by the identity of dissolved solute particles. (ex: vapor pressure lowering, boiling point elevation,
freezing point depression, and osmotic pressure)
 Boiling Point elevation:: Change in temperature = Kbm
 Freezing Point Depression: Change in temperature = Kfm
Chapter 15: Energy and Chemical Change:
 Energy: the ability to do work and to cause change. The Law of conservation of energy states that in any
chemical reaction or physical process, energy can be converted to one form or another, but it is neither
created nor destroyed.
 Chemical Potential Energy: energy that is stored in a substance because of its composition.
 **Know the conversion factors and units of energy listed on page 518 in table 15.1
 Specific Heat: the amount of heat required to raise the temperature of one gram of the substance by one
degree Celsius. The specific heat of liquid water is 4.184.
 Equation for heat: Q = mc(delta T)
 Universe = surroundings + system
 Enthalpy is the heat content of a system at a constant pressure. Hess’ Law states that if you can add two
or more thermochemical equations to produce a final equation for a reaction, then the sum of the
enthalpy changes for the individual reactions is the sum of the enthalpy changes of the individual
reactions. (**just think about changing the equations around like a puzzle, trying to get to the main
equation)