Chemistry I Review
The following guide is intended to be a study aid for the first exam in AP Chemistry in September. The exam will cover many of the topics listed below, although any information you learned in Chem I is potentially on the test. I advise that if you are unsure of any topic below, look it up over the summer and please ask during our review time when we return in September. The chapters below refer to the chapters in the AP Chem text book (Chemistry, Zumdahl & Zumdahl, 9th edition). Optional topics are noted in italics in the tables and lists below. You can use the internet sources such as Google, Youtube.com, Sciencegeek.net/chemistry, Chemtutor.com, ptable.com, Khanacademy.org to help you with these topics.
Good luck with your studies!
Know the definitions of:
Chemistry/Science Basics (Ch. 1)Scientific Method / Substance
Hypothesis / Matter
Experiment / Element
Theory / Compound
Law / Physical Property
Independent Variable / Extensive Property
Dependent Variable / Intensive Property
Scientific Notation / Chemical Property
Significant Figures / States of Matter
Accuracy / Solid
Precision / Liquid
Certainty / Gas
Unit Analysis / Vapor
Conversion Factor / Physical Change
Law of Definite Proportions / Chemical Change
Density / Law of Conservation of Mass
The Gas Laws (Ch. 5.1 – 5.5)
Pressure / Combined Gas Law
Barometer / Ideal Gas Law (PV = nRT)
Vapor Pressure / Ideal Gas Constant (R)
STP / Gas Stoichiometry
Boyle’s Law / Avogadro’s Principle
Charles’s Law / Dalton’s Law of Partial Pressure
Gay-Lussac’s Law / Molar volume of a gas @STP
Know also:
· Components of the scientific method: observation, hypothesis, experiment, conclusion.
· Determine precision and certainty (Percent error); I’ll give you the formula of
· Apply the law of Conservation of Mass
· Apply the law of Definite Proportions
· Density is mass/volume
· Use the Gas Laws (Boyle, Charles, Gay-Lussac, Combined)
· Determine the number of moles (or mass) of gas given mass, T, & P
· Ideal Gas Law: PV = nRT; how to work with this formula
· Determine density or molar mass of a gas using M=dRT/P
· Determine the partial pressure of a gas given its mole fraction and vice-versa.
· Convert units of pressure and temperature.
· Apply the law of conservation of mass
· How to add, subtract, multiply, divide numbers and get proper number of significant figures.
· Dimensional analysis to solve problems.
Know the Definitions of:
Atomic Nucleus (Ch. 2)Dalton’s Atomic Theory / Atomic Number
Nucleus / Isotope
Proton / Mass Number
Neutron / Atomic Mass Unit (amu)
Electron / Atomic Mass
Electrons in Atoms (Ch. 7)
Electromagnetic Radiation / Quantum Mechanical Model of atom
Wavelength (λ) / Atomic Orbital
Frequency (ν) / Principal Quantum Number
Amplitude / Principal Energy Level
Electromagnetic Spectrum / Energy Sublevel
Quantum / Electron Configuration
Planck’s Constant / Aufbau Principle
Photoelectric Effect / Pauli Exclusion Principle
Photon / Hund’s Rule
Atomic Emission Spectrum / Valence Electron
Ground State / Electron (Lewis) Dot Structure
Know also:
· Dalton’s Atomic Theory
· Charge on a Proton, Neutron, and Electron
· What makes up the nucleus
· Rutherford’s Experiment and what that showed about the nucleus
· Atomic Number = Number of Protons = Number of Electrons for a neutral element
· Mass number = Number of Protons + Number of Neutrons; know how to find any one given the other two.
· What shorthand notation means. i.e. C-14 means Carbon has 14 protons + neutrons, Carbon (C) is element #6 so it has 6 protons. C-14 then has 8 neutrons.
· Know how to calculate the atomic mass of an element given the mass and the abundance of the isotopes.
· Photons are bundles of energy and the units are Joules (J)
· Calculate the energy of an electromagnetic wave (E=hν)
· The atomic orbitals (s, p, d, f)
· How to fill the atomic orbitals of an element (Aufbau)
· Number of orbitals in an energy level, n. (# orbitals = n2)
· Number of electrons in an energy level, n (#electrons = 2n2)
· Electron orbital notation. Ex. B = 1s22s22p1
· Orbital Box Diagram. Example: Carbon
o 1s 2s 2p
· How to spot violations of Aufbau Principle, Hund’s rule, and the Pauli Exclusion Principle
Know Definitions of:
The Periodic Table & Periodic Trends (Ch. 7)Periodic Law / Transition element/metal
Group, Family, Column / Metal
Period / Alkali Metal
Representative Element / Alkaline Earth Metal
Inner Transition Metal / Nonmetal
Halogen / Noble Gas
Metalloid / Ion
Ionization Energy / Octet Rule
Electronegativity / Z*(Zeff) = Z – S
Know also:
· The classification of an element (metal, nonmetal, metalloid)
· How to locate an element by period, group, and block
· Whether the element is a representative element or a transition element
· Trends of Z* (Zeffective), atomic radius, ionic radius, ionization energy, electronegativity, and metallic character as you move across and down the periodic table.
· Size relationship between a neutral atom its corresponding ion
Know the definitions of:
Ionic Compounds & Metals (Ch. 2 & Chem I notes)Binary compound / Oxidation number
Chemical Bond / Polyatomic ion
Cation / Electron sea model
Anion / Delocalized electrons
Ionic bond / Metallic bond
Lattice Energy / Alloy
Formula unit / Interstitial alloy
Monoatomic Ion / Substitutional alloy
Hydrate
Know how to:
· Make a compound neutral when you have positive and negative ions.
· Show that column (group) of the element effects its oxidation number.
· Explain that lattice energy is affected by ion size and charge
· Identify elemental ions like oxide (O2-) and polyatomic ions like carbonate (CO32-)
· Write out a formula from the name of the compound (Ex. potassium nitrate is KNO3), and what to do when you have a transition element (Ex. Cobalt(II) chloride is CoCl2) and a hydrated compound (Ex. Barium iodide dihydrate is BaI2•2H2O).
· Write out a name when you have a formula. Ex. PbSO4 is lead(II) sulfate. Same ideas as previous bullet point, just in reverse.
· Note: You do NOT have to know the nomenclature rules for the oxyanions
Know the definitions of:
Chemical Bonding (Ch. 8)Covalent Bond / Structural formula
Molecule / Polar covalent bond
Lewis Structure / Electronegativity
Sigma (σ) bond / Bond dissociation energy
Pi (π) bond / Lone pair electrons
Endothermic / Allotrope (?)
Exothermic / VSEPR Model
Electron Domain Geometry (EDG) / Molecular Geometry (MG)
Hybridization (sp3, sp2, sp) / Linear
Tetrahedral/Trigonal Pyramidal/Bent / Trigonal Planar/Bent
Trigonal bipyramidal/See-Saw/T-shaped / Octahedral/Square Pyramidal/Square planar
Electron Geometry / Molecular Geometry
Know how to:
· Write the formula for covalent compounds (e.g. carbon dioxide)
· Determine type of bond between the two elements based on their electronegativity difference
· Qualitatively determine how number of bonds and bond length affects bond strength (i.e. double bonds are shorter and take more energy to break them than single bonds)
· Draw Lewis structure of molecules and polyatomic ions and be able to label bonding electrons and lone pair electrons.
· Be able to determine electron domain geometry (EDG), molecular geometry (MG), and bond angles of a Lewis structure.
· Know how having lone pairs and multiple bonds affects the EDG and MG
Know the Definitions of:
Chemical Reactions (Ch. 3)Chemical reaction / Types of Reaction
Reactant / Synthesis
Product / Decomposition reaction
Chemical equation / Single-Replacement reaction
Coefficient / Double-Replacement reaction
Precipitate / Combustion Reaction
Solute / Complete ionic equation
Solvent / Net ionic equation
Aqueous Solution / Spectator ion
Chemical Equilibrium (Ch. 13.1, 13.7 & Notes)*
Le Chatelier’s Law / Equilibrium
Reversible Reaction
* This material will be covered in greater detail in Ch. 13, which we will formally cover later in the year. The summary here are topics you should know from Chem I.
Know how to:
· Balance simple reactions
· Identify simple reactions as decomposition, synthesis, single replacement, double replacement or combustion.
· Recall the seven diatomic elements (H2, N2, O2, F2, Cl2, Br2, I2)
· Decomposition of carbonates yields a metal oxide and carbon dioxide (eg. CaCO3 D CaO(s) + CO2(g))
· Decomposition of chlorates yields a metal chloride and oxygen gas (eg. 2Al(ClO3)3 à 2AlCl3(s) + 9O2(g))
· Identify the physical states of reactants and products (s= solid, aq=aqueous, l=liquid, g=gas)
· Write the reaction using formulas if you are given the names and states of the compounds
· Write the net ionic equation given the complete ionic equation
· Apply Le Chatelier’s law, e.g.
CaCO3(s) D CaO(s) + CO2(g)
If we remove CO2 from a vessel, the reaction will shift to the right to create more CO2. If we add CO2 the reaction will shift left.
Know the Definitions of:
The Mole (Ch. 3)Mole / Percent Composition
Avogadro’s Constant / Empirical formula
Molar mass / Molecular formula
Hydrate
Know how to:
· Determine the molar mass of a substance (element or compound).
· Change grams to moles to number of atoms, molecules or formula units and back
· Determine the empirical formula of a compound given the elemental mass percentages (or masses) of the compound
· Determine the molecular formula of a compound given the elemental mass percentages (or masses) and the molar mass
· Determine the empirical formula of a hydrated compound given the mass percentages of salt and water in the compound
Know the Definitions of:
Stoichiometry (Ch. 3)Stoichiometry / Excess reactant
Mole ratio / Theoretical yield
Limiting reactant / Actual yield
Mol/mol; mol/mass; mass/mass conversion factors / Percent yield
Know how to:
· Perform mol-mol, mol-mass, mass-mol, and mass-mass stoichiometry calculations
· Determine the limiting reactant in a reaction
· Determine the amount of product(s) (mass or moles) formed in a limiting reactant problem
· Determine the mass (or moles) of excess reactant leftover in a limiting reactant problem
· Use gas volumes like moles in a gas stoichiometry problem
· Determine the volume of gas produced in a reaction at STP or other conditions
· Determine the percentage yield, theoretical yield, and actual yield of a reaction
Know the Definitions of:
Solutions (Ch. 11.1 & Notes)Molarity / Solution
Mass percent: %(m/m) & %(m/v) / Soluble/Insoluble
Saturated/Unsaturated Solutions / Miscible/Immiscible
Concentration (of a solution) / Dissolve
Parts per Million (ppm) / Mole Fraction
Colligative Properties / Freezing point depression
molality / Boiling point elevation
Kf and Kb / Heating curve/Cooling curve
Vapor pressure / Phase diagrams
Acids & Bases (Ch. 4.8, 14.3, 14.4, & Notes)*
Acid / pH
Base / pOH
Neutralization / [H+]
Strong Electrolyte/weak electrolyte / [OH-]
Strong acid/weak acid / Strong base/weak base
* This material will be covered in greater detail in Ch. 14 and 15, which we will formally cover later in the year. The summary here are topics you may know from Chem I.
Solutions
Know how to:
· Calculate the %(m/m) of a solution
· Calculate the %(m/v) of a solution
· Calculate the mole fraction of a solution
· Calculate the ppm concentration
· Calculate the molarity of a solution
· Convert between %(m/v) and molarity of a solution
· Prepare a solution of a certain molarity and volume (i.e. how many grams of solute to weigh)
· Determine the new concentration %(m/v) or molarity of a solution if it is diluted (C1V1=C2V2 or M1V1=M2V2)
· Determine the molality of a solution
· Determine the new boiling point or freezing point of a solution with a nonvolatile solute (∆T=Kb*m*i or ∆T=Kf*m*i)
· Use colligative properties to find the molar mass of a solute
· Read a phase diagram
· Calculate how much energy is absorbed or released upon heating or cooling a substance through a phase change
Acids and Bases
Know how to:
· Determine pH from [H+] or [OH-]; pH = -log[H+], pOH = -log[OH-]
· Determine [H+] from pH ([H+] = 10-pH) or [OH-] from pOH ([OH-]=10-pOH)
· Determine an unknown molarity by titration (neutralization)
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