Unit 1: Stoichiometry (CH 3 of textbook)

Know how to do all conversions types. Mass to moles, moles to particles, and molar ratios. Empirical formulas: change percent to mass, convert mass to moles, divide by smallest number

of moles to get molar ratio, round to nearest whole number, those whole number become the subscripts of empirical formula or the coefficient for a hydrate.

Molecular formula mass/empirical formula mass = mass ratio. Use mass ratio to multiply empirical formula subscripts, this results in the molecular formula.

Mass percent = mass of one substance/total mass Limiting reactants to determine theoretical yield Actual/theoretical x 100 = % yield Percent Error = |A-T|/T x 100 Gravimetric analysis (Filtration problems)

Unit 2: Electron structure and periodic properties (CH 6 and 7 of textbook)

1s2s2p3s3p4s3d4p5s4d5p6s4f5d6p7s5f6d7p Orbits hold 2 electrons that spin in opposite directions

o An s subshell contains 1 orbit = total 2 electronso A p subshell contains 3 orbits = total of 6 electronso A d subshell contains 5 orbits = total of 10 electronso An f subshell contains 7 orbits = total of 14 electrons

Valence electrons are electrons in the outermost energy level Know how to do light equations using speed of light and how to use the energy equation for

light wavelengths. Core electrons are all other electrons

o Example Br 1s22s22p63s23p64s 2 3d104p 5 Only outermost electrons are valence electrons

Effective nuclear charge the pull from the nucleus felt by the valence electronso Atomic number – core electrons = ENC

Screening electrons- core electrons that block the pull of the protons on the valence electrons Atomic radius increases to the left on the periodic table because the ENC on the valence

electrons is not as strong Atomic radius increases down a group because you are adding energy levels of electrons Ionization energy – energy required to remove an electron Ionization energy increases up a group because the electrons are closer to the nucleus due to

fewer energy shells Ionization energy increases to the right because of a larger ENC on the valence electrons which

pulls the electrons tighter to the nucleus.o Exceptions: Al compared to Mg because of full subshell on Mg, O is lower than N

because N has a half full subshell which result in a more stable energy level.

Electronegativity increases up and to the right, not including noble gases! Cations are smaller than their neutral atom fewer electron shells but the same number of

protons Anions are larger than their neutral atoms because they have the same number of protons but

more electrons (smaller effective nuclear charge on the valence electrons.)

Unit 3: Bonding, Geometry, and IMF’s (CH 8, 9, and 11 of textbook)

Ionic, covalent, and metallic bonding Ionic compounds- transferred electrons and lattice energy

o Lattice Energy – energy required to completely separate one mole of a solid ionic compound into its gaseous ions.

Covalent compounds – Shared electrons, molecular shapes, Lewis Structures Drawing Lewis structures

o Central atom is least electronegative atom Except H

Count up total number of electrons, single bond atoms to central atoms (each bond uses 2 electrons), fill the octet of outer atoms, fill octet of central atom. If you run out of electrons before octet is full, double or triple bond atoms.

Lewis structures formal chargeo Expected valence electrons - # of drawn valence electrons

Exceptions to the octet ruleo Under the octet: Boron makes only 3 bondso Expanded octet: Nonmetals that have a larger atomic number that Phosphorus

These nonmetals will have enough bonds and lone pairs to keep their formal charge at zero

Resonance structures – When electrons or bonds can be moved around a Lewis structureo Example: Ozone (O3) O=O-O or O-O=O

Shapes

2 electron domains

3 electron domains

4 electron domains

5 electron domains

6 electron domains

0 lone pairs Linear Trigonal planar Tetrahedral Trigonal bipyramidal

octahedral

1 lone pair Linear Bent Trigonal pyramidal

Seesaw Square pyramidal

2 lone pairs --------------- Linear Bent T-shaped Square planarBond Angles 180o 120o 109.5o 120o & 180o 90o & 180o

Hybridizaitono Sp3 hybridization = 1 s orbit and 3 p orbits that have hybridized together

Like a beefalo! These are formed from molecules that have 4 electron domains

o Sp2 hybridization = 1 s orbit and 2 p orbits hybridized with one unhybridized p orbit Formed from a molecule with 3 electron domains

o Sp hybridization = 1 s orbit and 1 p orbit hybridized with two unhybridized p orbits Formed from a molecule with 2 electron domains

Sigma bonds (σ)- every bond contains one of these Pi bonds (π)- A double bond contains one pi bond and one sigma bond

o A triple bond contains two pi bonds and one sigma bond Bond length from shortest to longest: triple, double, single

o Bond length is measured from the nucleus of one atom to the nucleus of the other bonding atom

Meaning smaller atoms will make shorter bonds than atoms with larger atomic radii

Bond dipoles occur when there is an electronegativity difference between two atomso The more electronegative atom has the partial negative end of the bond because it pulls

the electrons toward it Bond dipoles can tell us if a molecule is polar

o A molecule that contains polar bonds (bond dipoles) and is unsymmetrical in shape will be considered a polar molecule.

o A molecule that contains polar bonds but is symmetrical in shape will be considered a non polar molecule.

o Molecules that do not contain polar bonds are non polar. A molecules polarity tells us about the type of intermolecular force that holds the molecules

togethero Non polar molecules only exhibit London Dispersion Forces

Attraction between an instantaneous partial positive and partial negative. These partial charges come from the dispersion of the electrons around the molecule.

o Polar molecules exhibit dipole-dipole forces Attraction between the partial positive and partial negative of polar molecules.

o Hydrogen bonding a special type of dipole interaction that occurs between molecules that have a hydrogen physically bonded to a N,O,F

Strongest force Hydrogen bonding, dipole=dipole, LDF’s are the weakest forceo ALL MOLECULES EXHIBIT DISPERSION FORCES!!!o Only polar molecules exhibit dipole-dipole forces and hydrogen bonding.

Stronger forces result in a higher boiling point Higher boiling points stem from a larger molar mass also Coulomb’s Law: Energy associated with intermolecular attractions

o Can refer to IMF’s or Ionic bonds

Types of solids:o The following flow chart helps determine the types of solid that forms

Alloyso Substitutional vs interstitial

Substitutional: Mixture of metals where the atomic radii are similar so the metal replaces the metal that was previously present in the metallic bond

Interstitial: When the radii are not similar in size so the smaller metal can be placed in the small gaps between the metallic bonds

Unit 4: Gases (CH 10)

Pressure - number and force of collisions with the walls of a container Know how to use these gas laws (on the equation sheet)

o Combined gas Law P1V 1T 1

= P2V 2T 2

o Ideal Gas Law PV=nRT Not on the equation sheet is

o Molar Mass of a gas = DRTP

Where D= Density, R= Gas constant, T= Temperature, and P= Pressure Make sure Temperature is ALWAYS in Kelvin when doing a gas conversion Remember that partial pressures are additive for a gas

o Dalton’s Law of Partial Pressures Kinetic Molecular Theory

o Understand how this affects pressure Example: As temperature increases gas particles move faster, since particles

move faster there are more collisions with the container walls which results in an increase in pressure.

Mole fractions = moles of one substance/total moles of mixtureo Mole fraction times total pressure will give the pressure of that single gas

A gas is formed when enough energy is absorbed to break the intermolecular attractions between molecules.

22.4L/mol @ STP!!! Only when at STP!

Unit 5: Solutions (CH 4 and 13)

Calculating Molarity (mol/L) Using Molarity as a conversion factor

o Since M is mol/L it can be used to convert from liters of solution to moles of solute Ionic compounds dissociate into their individual ions when they dissolve

o AlCl3 splits into 1 aluminum ion and 3 chloride ionso This molar ratio can be used to determine concentration of a specific ion in a solution

Electrolyteso Compounds that fully dissociate are strong electrolytes

Ionic compounds and strong acidso Compounds that dissolve but do not fully dissociate are weak electrolytes

Weak acids and bases

o Compounds that do not dissolve or dissociate are non electrolytes Covalent compounds

Net ionic equations and spectator ions Molality = moles of solute/ kg of solvent Molality is used to calculate freezing point depression and boiling point elevation The solution process

o Know the difference between unsaturated, saturated, and supersaturated How do you supersaturate a solution?

o Like dissolves like This is not an explanation for why things dissolve, only a way to remember what

is soluble in what solvento Know how the intermolecular forces break apart solids

Positives are attracted to negatives Solvent molecules completely surround solvent particles Be able to draw and show the attraction between solvent particles and solute

particles

Unit 6: Kinetics (CH 14)

Remember that kinetics is the collision chapter!!!! Factors that affect reaction rates

o Temperature, concentration, physical state, and catalysts Remember that a rate is a measure of concentration per time

o M/s, M/hr, atm/s, atm/hro Concentration units can be Molarity for a solution or pressure units for a gas

Stoichiometry with reaction rates

aA + bB cC + dD

Rate = -−1a Δ [A ]Δt

= −1b Δ [B ]Δt

= 1c Δ [C ]Δt

= 1dΔ [D ]Δt

o Remember that reactants are disappearing so the rate is negativeo Products are appearing so the rate is positive

Rate Lawso Rate = k [A]m[B]n

Where m is the order with respect to A Where n is the order with respect to B

o To determine order of a reactant you can calculate it by doing:

( [A ]∈trial 2[A ]∈trial 1 )m =

rate∈trial2rate∈trial1

This will only work if the concentration of A changes but B stays the same Reactant orders are added together to give overall reaction order

Order Units for K Integrated rate law straight line graph0 order M/s [A] vs time1st order 1/s or s-1 ln[A]t = -kt + ln[A]o ln[A] vs time2nd order 1/ M*s or M-1s-1 1/[A]t = kt + 1/[A]o 1/[A] vs time3rd order 1/M2*s or M-2s-1

4th order M-3s-1

Rate of a reaction is dependent on the activation energy barriero Activation Energy Barrier- Amount of energy needed from the collisions of molecules for

a reaction to proceed.

o Activation energy graph showing an uncatalyzed reaction vs. a catalyzed reactiono Notice that a catalyst only speeds up a reaction by lowering the activation energy barriero The Ea is the activation energy

To calculate the activation energy we use the Arrhenius equation A derived Arrhenius equation gives us the following linear line.

lnk = −EaR

∗1

T+lnA

o The ΔE is the overall change in energy In this reaction the ΔE is negative so it is an exothermic reaction Products lower than reactants = exothermic Products higher than reactants = endothermic

Reaction Mechanismso The possible steps to a reaction o These steps must add up to produce the overall chemical equationo A catalyst is present in the beginning of the reaction and in the endo An intermediate is produced in one step then later consumedo The reactants leading up to the slow step or “rate determining step” will be present in

the rate law Look at pages 582-588 for help with reaction mechanisms

Unit 7: Chemical Equilibrium (CH 15 and 17.4)

Equilibrium is when the rate of the forward reaction is equal to the rate of the reverse reaction Q is the reaction quotient

o Calculate Q just like calculating K, products over reactants, raised to their coefficientso However Q calculations are at a given instantaneous moment, not necessarily at

equilibrium When Q > K the reaction will reverse back to produce more reactants When Q < K the reaction will proceed to produce more products When Q = K then the reaction is at equilibrium

Calculating Ko ICE, ICE, Baby!

Use an ice table to help organize data and thoughts of an equilibrium equation If change in concentration is not given use X Factor in coefficients to the ice table Change is negligible when k is less than or equal to x 10-5

o Plug values into equilibrium expression Products over reactants = k

Don’t forget coefficients in the equation turn into exponents in the expressiono K<1 means reactants are favored, k = 1 means even mixture of products and reactants,

k> 1 means products are favored Le Châtelier’s Principle

o Knowing how changes to an equilibrium system will cause a shifto Concentration

Comparing Q to K If Q < K shift to products If Q > K shift to reactants [Reactants] is increased equilibrium shifts towards products [Reactants] is decreased equilibrium shifts towards reactants [Products] is increased equilibrium shifts towards reactants [Products] is decreased equilibrium shifts towards products Due to an increase in collisions! Also this changes the value of Q which will shift concentrations until Q=k

o Temperature Increase in temperature will shift towards products for endothermic reactions Increase in temperature will shift towards reactants for exothermic reactions Decrease in temperature will shift towards reactants for endothermic reactions Decrease in temperature will shift towards products for exothermic reactions This is the only change that will change the value of k

o Volume and Pressure Increasing pressure will shift towards the side with fewer moles of gas Increasing volume causes a decrease in pressure, so it will shift towards the side

with more moles of gas Decreasing pressure will shift towards the side with more moles of gas Decreasing volume causes an increase in pressure, so it will shift towards the

side with less moles of gas Ksp equilibrium

o Since reactant is a solid in this case only products are part of expressiono Example PbCl2

Ksp = [Pb2+][Cl-]2

Ksp = [x][2x]2

If given molar solubility this plugs in for x

Unit 8: Acids, bases, and buffer solutions (CH 16 and 17)

Acid- Proton donor Base-proton acceptor Conjugate acid- base after it accepted a proton, ready to donate a proton in the reverse reaction Conjugate base- acid after it donated a proton, ready to accept a proton in the reverse reaction HCl + NH3 NH4

+ + Cl-

Acid Base CA CB Equilibrium expression

o Products over reactantso HA + H2O H3O+ + A-

Ka= ¿¿o B + H2O HB++OH-

Kb= ¿¿o ICE tables to determine concentration of reactants or productso When looking for the pH of a substance make sure to calculate the [H+] or [OH-] at

equilibriumo Smaller k values mean a weaker acid/base

Useful equationso pH = -log[H3O+]

[H3O+] = 10-(pH)

o pOH = -log[OH-] [OH-] = 10-(pOH)

o 14 = pH + pOH Titrations

o Equivalence point - the point in a titration when the moles of acid = moles of baseo Strong acid & Strong Base

Both substances fully dissociate Initial pH = -log[Strong Acid] Or initial pH = 14 – (-log[Strong Base]) Before equivalence point

Calculate the amount of mole of acid or base left over Divide by total volume to get Concentration Calculate pH based on new concentration of acid/base

At equivalence point

pH = 7 After equivalence point

Calculate the number of moles of acid/base present in excess Divide by total volume to get concentration Calculate pH based on new concentration of acid/base

o Strong acid/base & weak acid/base Only one substance fully dissociates Initial pH is based on ice table and k value Before equivalence point

Calculate number of moles of acid/base left over and moles of the conjugate that has been produced

Plug into the David Hasselhoff equation (Henderson-Hasselbach)

pH = pKa + logConjugatebase

acid At equivalence point

Calculate moles of conjugate base that are produced Divide by total volume to get concentration Write a dissociation equation for the conjugate base and water Calculate the Kb usind Kw = Ka x Kb

Use an ice table and the new Kb value to calculate [OH-] 14-pOH= pH

½ way to equivalence point the pH = pKao Know what a titration curve would look like for these instances

pH of salt solutionso Cations act like an acid

Only cations that would be part of a weak base affect the pHo Anions act like bases

Only anions that are the conjugate of a weak acid affect the pH 7 Strong acids

o HCl, HBr, HI, HClO4, HClO3, H2SO4, HNO3

Strong baseso Alkali metals bonded to OH-

o Ca, Sr, Ba bonded to OH-

o Strong acids and strong bases fully dissociate into their ions

Unit 9: Thermodynamics (CH 5 & CH 19)

Endo vs exothermico – ΔH is exothermico + ΔH is endothermic

Enthalpy refers to the energy transfer of a given processo Enthalpy of reaction or ΔHrxn

ΔHrxn = Σ ΔHproducts - Σ ΔHreactants

Remember to account for stoichiometric coefficients Unit for enthalpy is in kJ/mol

o Enthalpy of formation or ΔHf

Values for the formation of a compound from its standard state elements ½ N2(g) + O2(g) NO2(g) ΔHf =33.84 kJ/mol Unit is in kJ/mol Elements at their standard state have an enthalpy of 0 These are the values used to find the enthalpy of reaction

o Hess’s Law Reactions will add up to give an overall reaction These enthalpies will add to give the enthalpy of the overall reaction Equations may need to be multiplied by a number Whatever is done to an equation must be done to the enthalpy If a reaction is reversed the enthalpy for that reaction is the opposite of the

forward reactiono Calorimetry

q=mc ΔT Where q is heat

Can also be thought of as enthalpy M is mass measured in grams

Mass of a solution needs to be converted to grams using the density of the solution

If density is not given use 1g/mL which is the density of water C is specific heat capacity

Measured in a unit of Jg∗K

Specific heat of water is 4.184 ΔT is change in temperature

Can be measure in K or oC Entropy- Measure of the disorder or “randomness” of particles

o Things tend to move towards disorder or chaoso So a + ΔS means disorder whereas a – ΔS means ordero Entropy increases as you move from solid to liquid to gaso Entropy will increase as a reaction proceeds towards more moles of a substanceo Entropy increases when a solid is dissolved to form a solutiono Entropy is labeled as ΔS

Calculating ΔS = Σ ΔSproducts - Σ ΔSreactants

Questions will ask to predict the sign of the entropy change or calculate the entropy of a process

ΔS= - ΔHsys/To Unit for entropy is in J/mol*K

Gibb’s free energyo True measurement of the thermodynamic favorability of a process taking into accound

ΔH, T, and ΔSo + ΔG means a process is not thermodynamically favoredo – ΔG means a process is thermodynamically favoredo ΔG can be calculated using ΔG = Σ ΔGproducts - Σ ΔGreactants

Unit will be in kJ/molo When given enthalpy and entropy use a different calculation

ΔG = ΔH-T ΔS Pay attention to units!!

This table is Very important but you don’t need to memorize it

Think about how you would calculate ΔG using enthalpy, entropy, and temperature

This will tell you if ΔG is positive or negative

o Free energy and the equilibrium constant

ΔG = -RTlnK

Use this equation to determine the equilibrium constant

Unit 10: Electrochemistry

Redox reactions

Reduction is the gain of electrons, the ox # gets reduced

Oxidation is the loss of electrons, the ox # goes more positive

Remember LEO GER or OIL RIG

Use Reduction potential to calculate the Ecell

o ECell = Ecathode - Eanode

o If Ecell > 0 means process is thermodynamically favored

o If Ecell < 0 means process is no thermodynamically favored

Cathode

o Reduction occurs at the cathode

o Substance with the largest reduction potential is the substance at the cathode

o Positive ions from solution gain electrons and stick to the cathode

Anode

o Oxidation occurs at the anode

o Substance with the smallest reduction potential is the substance at the anode

o Electron are lost from a neutral atom the atom forms a positive ion which goes into solution

Electrons are transferred from anode to cathode

ΔG = -nFEcell

o F is Faraday’s constant

o N is number of electrons

o Pay attention to units!!!

Salt bridge

o Na+ and NO3-

o Anion in salt bridge moves toward the anode

o Cation in salt bridge moves toward cathode

Test Taking Tips!!!!

Multiple Choice

Don’t leave questions unanswered

Skip questions you don’t know and come back to them

Circle Questions and put a line through it if you have no idea!

o Then guess on these questions before you submit

Rule out answers you know are wrong

Pay attention to units

Minor details can be the difference in a right or wrong answer

Free Response

Be clear and concise with your answer

Box your correct answer

Include proper units always

SHOW YOUR WORK

o Partial credit is sometimes given for showing work

Periodic Trends are NOT an explanation or a reason

o You must explain your answer based upon the reason why a trend occurs

Convert to moles if you are stuck.

Don’t spend too much time thinking about one part. You need to answer as many questions as possible!