AP Chemistry Review Notes

Unit 1: Atomic Structures and Properties

Mass Spectrometry:

• Shows the mass of the isotope type of an element

• Relative abundance adds up to 100%

• From this, you can determine the amu by multiplying amu by percent, adding them all together, and dividing by 100

Mole Calculations:

• particles * avogadro's number (6.0231023) = moles

• grams * molar mass = moles

Percent Mass:

• Molar mass of certain element / molar mass of whole molecule * 100% = mass percent

Identifying Pure Substances:

• If mass percent is higher than it should be, an impurity contains that element

• The same is true for that in reverse

  • Adding extra mass from impurity without that element lowers the percent

• Empirical Formula: 

  • Divide my MM to get moles, then take the lowest number of moles and divide everything by that to get whole numbers

Gravimetric Analysis:

• Using gravity to separate two mixtures

  • ex. sand and water put through filter

  • Often part of a multi-step process

• Used in labs

Electron Configuration:

• 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6 ect…

• Can also use the closest noble gas before the element, then continue the configuration

  • ex. [Ar] 4s^2 3d^2 = Ti

• d block: subtract 1

• f block: subtract 2

• For the boxes, fill in all the up arrows first, then down
  

Periodic Trends:

• Coulomb's law: distance and charge explain all of them

  • Atomic radius: down and to the left

    • Radius of an atom.

  • Ionization energy: up and to the right

    • Energy to remove one electron

    • First ionization energy is the ionization energy from the furthest ring of electrons from the nucleus, the easiest to remove

  • Electronegativity: up and to the right

    • Pull atom has on its electrons

• Only thing you need to know for charge is the # of protons in nucleus

Formulas to Solve for Energy and Wavelength:

• E = hv

• c = 𝝺v

• E = hc/𝝺
  

Photoelectron Spectroscopy (PES) Diagrams:

• Photoelectron Spectroscopy (PES) diagrams can be used to determine elements

  • By seeing the number of electrons

• Relates binding energies with electron orbitals

• Closest to the left is closest to nucleus

• Smaller atomic radius means more attraction, so all the peaks are closer to the nucleus on the left of the graph

  • Due to charge, Coulomb's Law again (it will come up a lot later as well)

Vibrational Properties:

• Bonds can be stretched symmetrically or asymmetrically, bent, vibrate, and stretch along the x or y axis
  

Chemical Reactivity:

• Most reactive metals are group 1 metals, most especially Francium

  • Left and down are most reactive

• Most reactive nonmetals are group 17 nonmetals

  • Right and up are most reactive

Bond Polarity:

• The greater the difference in electronegativity, the more polar the bond is

Unit 2: Molecular and Ionic Compound Structures and Properties

Bonding Properties:

• Like dissolves like, polar dissolves in polar (water)

  • Nonpolar only dissolves in nonpolar solvents
    

Metallic Solids Characteristics:

• Substitutional alloy

  • Replace original metal with a new one

  • Stronger and malleable

• Interstitial alloy

  • Fits in-between structure

  • Brittle and unbendable

Metallic Properties:

• Metals are surrounded by a sea of electrons, so the electrons can carry charge through the molecule in between gaps between atoms

• Malleable because the sea of electrons can be compressed and moved to allow for the movement of atoms
  

Ionic vs Covalent:

Lewis Dot Structures Steps:

• Steps to draw Lewis Dot Structures

  • Count VE

  • Draw structure with single bonds connecting to middle

  • Recalculate VE

  • Give to outside first

  • Recalculate VE

    • If leftover VE give to middle

    • If no left over, double/triple bond

VSEPR Theory:

• 2 bonds, 0 one pair: Linear 180 sp

• 3 bonds, 0 lone pair: Trigonal Planar 120 sp^2

• 2 bonds, 1 lone pair: Bent 120  sp^2

• 4 bonds, 0 lone pair: Tetrahedral 109.5 sp^3

• 3 bonds, 1 lone pair: Trigonal Pyramid 109.5 sp^3

• 2 bonds, 2 lone pair: Bent 109.5 sp^3

Unit 3: Intermolecular Forces and Properties

Electrostatic Forces:

• Exist between molecules and atoms within molecules

• Between molecules and called IMF: have three different types

  • London Dispersion: all have it, nonpolar have only it

  • Dipole-Dipole: polar molecules

  • Hydrogen Bonds: H directly bonded to F, O, or N

• Strong IMF have:

  • High BP, high MP, high viscosity, high surface tension, low vapor pressure!

Physical Change vs Chemical Change:

• Physical Change

  • Relate to IMF changes

    • Because of this, the highest IMF means highest boiling point

    • More electrons -> more polar -> more IMF

  • ex. phase changes

• Chemical Change

  • Bonds are broken and formed to create a new substance

Hydrogen Bonding:

• Because N, O, and F are most electronegative, they pill electrons away from the H they are bonded to and take on a partial negative charge, and make the H have a partial positive charge that attracts to another partial negative charge
  

Chromatography:

• Mased of the mobile phase, which is the liquid used

  • For example if you use polar water, the substances that move with the water will also be polar

  • Based off IMF

Distillation:

• Based off differences in boiling points

• You can heat up a substance between two BPs, and you can collect the liquid that turns into a gas and separate the two liquids

Solids vs Liquids vs Gasses:

• Solids vibrate

• Liquids move around the bottom

• Gasses move quick everywhere, further apart
  

Kinetic Molecular Theory:

• There are some deviations from ideal behavior

  • Low temperature and high pressure (low volume)

  • This will cause more deviations

  • Also, the ones with strongest IMF deviate most

Properties of Gas:

• Pressure and volume are inversely proportional

• Volume goes up, temperature goes up

• More moles of gas, greater volume
  

Ideal Gas Law:

• PV = nRT

  • P = pressure

  • V = volume

  • n = moles of gas

  • R = .08206

  • T = temperature

• n = m/M

  • Moles = mass / molar mass

• PV = mass * RT / MM

  • Deviated from above equations

  • MM = mass * RT / PV

  • More deviations can be derived from the equations

  • Less mass moves faster at same temperature
    

Dissolving:

• Positive charged ion attracts the O- of the H2O

• Negative charged ion attracts the H+ of the H2O
  

Molarity and Particle Views:

• molarity = moles/ liters

• Rank the six solutions above in order of increasing molarity. Pay attention to volume, and some have equal concentration

  • C,D, and E (tied); A and F (tied); most concentrated is B

Coulomb’s Law and Solubility:

• Ionic compounds can dissolve in polar liquids like water because the ions are attracted to either the positive or negative part of the molecule

• The smaller the ions, the closer together they are, and the harder it is for the water molecules to pull the ions away from each other. The greater the charge of the ions, the harder it is for the water to pull them away as well.

Beer-Lambert Law:

• Measures the concentration of colored solutions

• A = abc

  • A = Absorbance

  • a = molar absorptivity

  • b = path length

  • c = concentration

• Things can cause deviations in absorbance

  • More water will cause less absorptivity

  • A fingerprint will cause more absorptivity

Unit 4: Chemical Reactions

Chemical Change:

• Breaking and forming of bonds

  • Synthesis

  • Decomposition

  • Single Displacement

  • Double Displacement

  • Combustion

  • Acid Base (Neutralization)

  • Oxidation-Reduction

  • Precipitation
    

Balanced Equations:

• SNAP ions,

  • If it has one of the four, it will be aqueous

  • If not, it will form a precipitate

Stoichiometry:

• Use mole ratio to convert between molecules

• Uses Law of Conservation of Mass

  • Whatever number of each element you have on one side, you have to have on the other

• Limiting Reactant

  • The reactant with the least moles, and it is what you divide everything by to get the empirical formula
    

Acid Base Reactions:

• Also called Neutralization reactions

• Strong acid + Strong base → salt + water

Redox Reactions:

• Reduction-Oxidation reactions

  • RED = reduction

  • OX = oxidation

• When an electron is transferred

• OIL RIG

  • OIL = Oxidation Is Loss (of electron, more positive)

  • RIG = Reduction Is Gain (of electron, more negative)

• Some charges are typically the same

  • Hydrogen +1 (typically)

  • Oxygen +2 (almost always)

  • Most Group 1 are +1

• Redox Titrations

  • Change color when there is a change in oxidation number

Unit 5: Kinetics

Lowering Reaction Rates:

• Introduce catalyst

  • Lowers activation energy

  • Only thing that changes activation energy

• Increase temperature

  • Makes more particles have enough energy to meet the activation energy

  • Particles move faster to create more collisions

• Increase concentration

  • More particles, more collisions

• Grinding up solid

  • Greater surface area

Rate Laws:

• Determined from concentration and rate changes

• Zeroth order

• When concentration is changed, the rate is not changed

• [A] over time

• First order

  • When concentration is doubled, so does the rate

  • Constant half-life

  • ln[A] over time

• Second order

  • When concentration is doubled, rate is tripled

  • 1[A] over time and goes up

Half-Lifes:

• How long it takes for half the sample to remain

• Constant half-lives only for First Order

Reaction Mechanisms:

• Rate = k [concentration] ^ # of moles

• The step that takes the most time is the slow step

  • This is the step that determines the rate of the reaction

Reaction Rate Graphs:

• # of peaks is the # of steps

• Largest peak is the slow step

• Exothermic ends lower

• Endothermic ends higher

Unit 6: Thermochemistry

Endothermic vs Exothermic:

• Endothermic

  • Thermometer colder

  • Taking heat from surroundings

• Exothermic

  • Thermometer hotter

  • Giving off heat to surroundings
    

Heat Transfer:

• q = mcΔT

• Specific heat capacity is how much energy needed to raise one gram one Kelvin
  

Conservation of Energy:

• Used to figure out the unknown when dissolved in water

• You can determine this from the heat change

  • Heat gained by water = heat lost by metal

Bond Energy, Length and Strength:

• Calculating ΔH

  • Bonds broken - Bonds formed

  • Products - Reactants

  • Hess's Law

    • Getting to target equation from combining other equations

    • ex. flip equation, multiply ΔH by -1

• Bond length

  • Single bonds are long and less strong

  • Double bonds are in the middle

  • Triple bonds are short and strongest

ΔH and ΔS:

• The ΔH is dependent on the IMF of the solute and solvent

• ΔS is entropy, which is disorder of a system

  • More entropy means more moles of gas, more pressure, higher temp, ect…

Conservation of Energy when Mixing:

• -Q (energy lost by system) = +Q (energy gained by surroundings)

  • Total energy always conserved

• For example :

  • When room temperature water T1 (system)  is mixed with cold water T2 (surroundings), the final temperature T3 will be in-between.

  • Q1 + Q 2 = 0 and energy is conserved
    

Calorimetry:

• Used to determine heat transferred

• Q=mcΔT

Unit 7: Equilibrium

Chemical Equilibrium:

• When the rate of the forward and reverse reactions are equal

  • Products are being made at the same rate and reactants

  • Not necessarily meaning there is the same amount of products and reactants

• K, the equilibrium constant, shows the proportion of reactants formed at a specific temperature that has achieved equilibrium
  

Manipulating Q and K:

• K (equilibrium constant) represents the ratio of products to reactants at equilibrium

  • When manipulating many K values to create a new solution

    • Flipped reaction: K-1

    • Multiplying reaction: Kx

    • Adding reactions: KK

• K₃ = 2.9

• Q (reaction progress) shows products to reactants at any point in a reaction

  • No solids or liquids are included in either of these, only aqueous and glasses

Q vs K:

• Q<K: reaction proceeds forward

• Q>K: reaction proceeds backward

• Q=K: reaction is at equilibrium

ICE Tables:

• One way to calculate K

  • Uses initial concentration, the change in concentration in forms of x, and the concentration at equilibrium

Magnitude of K:

• Can be used to determine relationship between reactants and products at equilibrium

  • Large K favors forward reaction (products)

  • Small K favors reverse reaction (reactants)

Le Chatelier’s Principle:

• Describes changes in systems to shift equilibrium

Common Ion Effect:

• The solubility of a hydroxide can be greatly increased with the addition of an acid

• In the same way, a salt’s solubility can be increased greatly by adding a common ion

  • ex. if MgCl₂ is added to a saturated Mg(OH)₂ solution, additional Mg(OH)₂ will precipitate out. The additional Mg²⁺ ions will shift the original equilibrium to the left, thus reducing the solubility of the magnesium hydroxide.

Unit 8: Acids and Bases

Bronsted-Lowery:

• Acids are proton donors

  • Acids create conjugate bases called anions

• Bases are proton acceptors

  • Bases conjugate acid now has a proton (Hydrogen ion)

K_W and Temperature:

• As T increases, the pH of pure water decreases

  • This is not because water is acidic

  • A solution is acidic if [H+]>[OH-]

Titrations:

• Allowed to be used to determine the concentration of the titrant and the pK_a for the weak acid and pK_b for a weak base

  • pK_a is the pH at the 12 equivalence point

Buffers:

• Buffers are able to resist pH changes when mixing a conjugate acid-base pair

• pH of buffers are determined by the pK_a of the weak acid

• Only effective if it has sufficient amounts of both members of a conjugate acid-base pair

Unit 9: Application of Thermochemistry

Entropy:

• Changes that result in more moles, higher temperature, greater volume, solid → liquid → gas, formation of more molecules

• Overall greater “chaos” and “disorder”

ΔG = ΔH - TΔS:

• ΔG = Gibbs Free Energy (kJ)

  • Negative ΔG means that the reaction is thermodynamically favored (spontaneous)

• ΔH = enthalpy change (kJ)

• T = temperature (Kelvin)

• ΔS = entropy change (kJ/K)

  • Often will have to change J to kJ