W e're no strangers to love

You know the rules and so do I

A full commitment's what I'm thinkin' of
You wouldn't get this from any other guy
I just wanna tell you how I'm feeling
Gotta make you understand
Never gonna give you up
Never gonna let you down
Never gonna run around and desert you
Never gonna make you cry
Never gonna say goodbye
Never gonna tell a lie and hurt you
We've known each other for so long
Your heart's been aching, but you're too shy to say it
Inside, we both know what's been going on
We know the game and we're gonna play it
And if you ask me how I'm feeling
Don't tell me you're too blind to see
Never gonna give you up
Never gonna let you down
Never gonna run around and desert you
Never gonna make you cry
Never gonna say goodbye
Never gonna tell a lie and hurt you
Never gonna give you up
Never gonna let you down
Never gonna run around and desert you
Never gonna make you cry
Never gonna say goodbye
Never gonna tell a lie and hurt you
 We've known each other for so long
Your heart's been aching, but you're too shy to say it
Inside, we both know what's been going on
We know the game and we're gonna play it
I just wanna tell you how I'm feeling
Gotta make you understand
Never gonna give you up
Never gonna let you down
Never gonna run around and desert you
Never gonna make you cry
Never gonna say goodbye
Never gonna tell a lie and hurt you
Never gonna give you up
Never gonna let you down
Never gonna run around and desert you
Never gonna make you cry
Never gonna say goodbye
Never gonna tell a lie and hurt you
Never gonna give you up
Never gonna let you down
Never gonna run around and desert you
Never gonna make you cry
Never gonna say goodbye
Never gonna tell a lie and hurt you
The FitnessGram™ Pacer Test is a multistage aerobic capacity test that progressively
gets more difficult as it continues. The 20 meter pacer test will begin in 30 seconds. Line up at the
start. The running speed starts slowly, but gets faster each minute after you hear this signal. A
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test is over. The test will begin on the word start. On your mark, get ready, start.
mklnkn kay Get that 5 Ffellahs
w
Also let me know if the chem final has changed and if there's any information that n
eeds to go here/leave
AP Chemistry Score Calculator–

2
It covers 60 multiple-choice questions (90 minutes total) and 7 free-response
questions (105 minutes total). Section I consists of the multiple-choice
questionsw13questions. These make up 50% of the exam score, and Section II
consists of 3 long-answer FRQs and 4 long-answer FRQs.i
Good ish link:https://blog.prepscholar
.com/ap-chemistry-exam
Use college board and the CED.
You can share:

● The impact you made

● Leadership & ienitiative you showed
● What you learned
● The role you played

Think there are 60 mc not 75

now there are 6o mcq but for
some reason they arbitrarily take 10
out before grading, so you solve

60, but get graded based on

50 of them?
OK. GO TAKE THE EXAM NOW. ENOUGH TALKING.

Periodic Table

Atom- Smallest form of matter. Over 99% of the mass is composed of protons and neutrons
in the nucleus, which is orbited by electrons.
Element- Specific configuration of an Atom
Compound- specific configuration of elements, whole number ratios

3
Horizontal Rows = Periods
Vertical Rows column? = Groups

Groups:
Group IA/1 - Alkali Metals
Group IIA/2- Alkaline Earth Metals
Group B/3 - Transition Metals
Group VII/17 - Halogens
Group VIII/18 - Noble Gases

Atomic Mass - Atomic Number = Number of Neutrons

Different Number of Neutrons are called Isotopes
Carbon-12 is normal
Carbon-14 is an isotope of 6 protons and 8 neutrons

yum

Isotope (Nuclide) Notation:

4
(26) Emre Kayra Kaya
1:48 AM Dec 27
Add: “w”

Anonymous

9:15 PM Mar 7
Delete: “w”

Anonymous

7:48 AM May 2
im not making it

Blake Chang

3:30 PM Today
Add space

5
Alexis Wingenbach

2:55 PM Today
Format: indent first line, indent left

William Minho Chung

2:55 AM Apr 23
Add: “there are 6o mcq but for some reason they arbitrarily take 10 out before grading, so you solve 60, b…”

Alexis Wingenbach

2:55 PM Today
Add tab (8 times)

6
Heba

10:59 AM Dec 5
Format image: position (2 times)

7
Ananya shivalika Kandaswamy saravanakumar

Nov 3, 2023
Delete image
Add image you guys are fucking niggers

Gina Lin

Sep 9, 2023
Add: “column?”

Arvindh Krishna

Sep 10, 2023

yeah its a column

Arvindh Krishna

Sep 10, 2023

Replace: “Number” with “Numbers”

Deepshikha Banerjee

7:50 PM Jan 21
Add: “(Nuclide)”

sophie

Aug 10, 2023

Add: “/”

Jungmin Lee

8
3:57 PM Apr 24
Delete: “Free Energy”
You're suggesting
11
10
9
8
7
6
5
4
3
2
1
1
2
3
4
5
6
7
8
9
10

Element X is normally X-98

Element X could be Isotopes in 92, 94, 96, etc.

Average Atomic Mass = ( Mass Isotope 1*%Abundance Isotope 1 ) + ( Mass Isotope 2*j %Abundance Isotope 2 ) +
( Mass Isotope 3*%Abundance Isotope 3 ) + etc…etcso sigma...

Valence electrons.

Number of Valence Electrons = Main group number - (charges)(electron number)

Electrons = -1 charge
Example: N 2+ = 5 (main group number) - 2 (2 charge electrons) = 3 valence
electrons

MOLES

Avogadro's Number : Defined constant for the amount of molecules present in one mole.
( 6.022 x10^23 )

Convert Moles to Grams (vice versa)

Moles = (gram/molar mass)

9
 Grams = (Moles x Molar Mass)

Moles of a gas are found with n=PV/RT ( Rearranged Ideal Gas Law. )
Most likely will occur at STP (Standard Temperature & Pressure)
Pressure = 1 atm
Temperature = 273 K (0 C)
At STP one mole of an ideal gas will be at 22.4 Liters.

P = Pressure in atm
V = Volume in liters
n = # of moles
R = Gas Constant ( .0821 L*atm/mol*K)
T = Temperature in Kelvin ( Take Celsius and add 273 to get Kelvin )(Test won’t give
Fahrenheit)

Moles = (molarity)(liters of solute)

Percent Composition- percent of mass of each element in a compound

Divide each element's molar mass by compounds molar mass
Ca(NO3)2 = 164.10
Calcium = 1 = 40.08 x 1 = 40.08

Nitrogen = 2 = 14.01 x 2 = 28.02

Oxygen = 6 = 16 x 6 = 96.00
Ca = 40.08/164.10 = 24.42%
N = 28.02/164.10 = 17.07%
O = 96/164.10 = 58.50%

Empirical Formula

Empirical formula can be found by calculating the moles of each element presented, and
then dividing all mole values by the smallest number found. The values given by the
division will be the empirical formula subscripts.

10
Electron Configurations

Nucleus is positive, pulling in electrons which are negative, closer they are stronger the
attraction is i.e More protons= more attraction, closer electrons= more attraction

Electrons are repelled by other electrons, an electron between a valence and nucleus
causes the valence to be weaker (called shielding)

Completed shells are stable, atoms strive for completed shells

Aufbau principle- when building electron configurations, electrons go into shells in order of
increasing energy

Pauli Exclusion Principle- two electrons in a shell will not spin the same way

Hund’s rule- when an electron is added it will take up the lowest energy orbital

Amount of energy from an electron is dependent on how far away it is from the nucleus

Bohr Model - Quantum model for the atom where electrons are arranged in specific orbits
around the atom - used this to predict the emission spectrum of hydrogen given the
electron distribution. Problem - good for predicting some emission spectrums but does not
work with most. Also, electrons do not “orbit” in a perfect, flat plane.

If an atom is exposed to electromagnetic levels greater than the ones exerted by the
electrons it holds the electrons can be ejected.

Electron Configuration w/ periodic table


Use Periods (horizontal) as first number, then letter of the group, then amount
Examples
Hydrogen: 1s^1
Helium: 1s^2
Lithium: 1s^2, 2s^1
Sodium: 1s^2, 2s^2 2p^6, 3s^1
You can also ‘shorthand’ by using the Noble gasses as starting points
Argon: [Ne] 3s^2, 3p^6

Moving left to right from periodic table atomic radius (distance from nucleus to valence
electrons) decreases (Zeff increases while valence electrons are added to same energy
level, so overall the electrons are pulled in more)

11
Moving down a group, atomic radius increases (more shells mean the valence electrons are
further out)
X

Removing electron creates a positively charged cation (X1+), the energy to remove the
first electron from a gaseous ion is the ionization energy, the energy to remove the next
electron is the second ionization energy, which is always higher than the first ionization
energy as the number of protons remains the same while there is one less electron, so
less force repelling and the same force attracting, overall requiring more force to remove
the electron.

If the electron being removed drops a shell, ionization energy increases MASSIVELY.

Electronegativity
Small atoms have larger electronegativity levels
Atoms close to having stable electron amounts have stronger electronegativity levels

BONDS

Bonds - sharing of electrons, more electrons shared = stronger bond

So single bonds (1 σ) have the least bond energy
Bond Energy Strength: triple (2 π and 1 σ)>double (1 π and 1 σ)>single (1 σ)

Ionic Bond- Bonds held by electrostatic forces, one will be + (cation(red)) while the other is
- (anion(black))
Ionic bonds are where the anion completely removes the electron from the cation
Strength of bond depends on size; smaller size=greater attraction

Metallic Bond- Metals in the bond are usually stationary and share a sea of electrons,
making them good conductors as they give up electrons easily,this also makes the metals
malleable

Metals can bond with each other and form alloys

Interstitial Alloy- metals with VERY different atomic radii combine
Example is carbon filling in the cracks around iron to make steel
Substitutional Alloy- Two metals with similar atomic radii combine
Example is Zinc and Copper, it is much more structured and uniform


12
Covalent Bonds- Two atoms share electrons and both count the electrons as part of their
valence shell (nonmetals)
The first covalent bond is called the sigma bond (σ), and has the lowest bond energy
and longest bond length, while the following bonds (2nd and 3rd) are called pi bonds

Network Covalent Bonds-these bonds are formed when a group of atoms share covalent
bonds into one big electron sharing cluster, it is very hard to break and has a high melting
point
Good contenders are Silicon and Carbon since they only have 4 valence electrons
(ex. SiO2) <--saw this as a suggestion from SiO4, its been years since i did chem which is
correct?

Polarity- Atoms usually will have an electronic imbalance of sorts due to uneven
distribution of electrons (very based on symmetry)
Dipole- The most electronegative element will pull on electrons but not enough to become
ionic, but will keep them on one side of the atom, H2O for instance has 2 H+ on one side,
connected to O- this makes the O- side negative, and the H+ side positive, more charges
means more dipole moment, which is how a dipole is measured
Dipole moments can attract each other, causing the H+’s in H2O to be attracted (they do
not form a bond) to the O-’s in other water molecules (this is a hydrogen bond described
below)

Intermolecular forces- forces that exist between 2 or more covalent molecules that hold
them together. (ex. London Dispersion forces, dipole/hydrogen bonds.) Stronger IMFs =
higher boiling point, lower vapor pressure

Hydrogen Bonds- Special type of dipole moment attraction that involves a positively
charged hydrogen to bond to a negative end of another molecule, usually one with high
electronegativity (H―F, H―O, H―N).
Extremely strong bonds, such as hydrogen, give up it's only electron so it uses all its
strength to hold onto it, molecules with these bonds have high boiling/melting points. A
group of H-F molecules, or any molecule w hydrogen bonding,are oriented so that the very
positive H of one molecule is attracted to the very electronegative F of another molecule.

London Dispersion forces- With large amounts of electrons moving in random patterns it is
possible for a majority to move to a specific side, creating a quick bit of polarity (Dipole
Moment), which can in turn disrupt other molecules. More electrons means more chance
of them congregating ( more polarizable ), so molecules with high boiling points are the
most vulnerable. Larger molecules are more polarizable so have more electrons and will
more likely have London dispersion forces
To compare the forces, compare the IMFs - larger is stronger

Bond order = (# of sigma and pi bonds/ # of atoms)

13
Formal Charge = Group Number - Lone electrons - # of sigma and pi bonds
(Group Number - dots - sticks = formal charge)

Bond Strength + Intermolecular Forces ( IMFs )

Strongest to weakest

Covalent Bonds
( Usually liquids. )

Network Covalent Bonds are strongest possible (think diamond).

(sharing valence electrons means you need to break 2 sets of valence electrons, making it
the toughest to break)
(Melting/boiling point is lower than ionic)
Use 4,5,6 bonds below to base on other covalent bonds strength
Electrons are shared.

Nonmetal-Nonmetal.

Ionic Bonds
Usually solid.
Energy to brekijjjjjak based on coulombic attraction
Bad conductors in solid form as electrons don’t move much.
Generally good conductors in aqueous solution.
Electrons are transferred.
Metal-Nonmetal.

Metallic Bonds
Transition metals seem to be strongest of all metals
High melting point, usually below ionic
Generally good conductors as a solid. ( ‘Mobile sea of electrons.’ )
Metal only.

Hydrogen Bonds ( Actually an IMF despite ‘bond’ being in the name. )

Special/specific kind of Dipole attraction
Most commonly between H and F,O,or N ( “Hydrogen bonding is FON!” )- corresponding
F,O, or N must have a pair of unpaired e-
Strongest IMF.
Contains both LDFs and DD.

Non-Hydrogen Dipole bonds ( Dipole-Dipole - DD)

Usually long lasting
Stronger than LDFs but weaker than HB.
Contains LDFs

14
Polar

London Dispersion Forces ( Sometimes called Induced Dipole )

Only creates a dipole temporarily
Weakest IMF.
Is the reason why larger molecules tend to have higher boiling points/lower vapor pressure
than smaller ones, can be used to help determine which molecules have the stronger IMFs
if they both have D-D and/or HB.
Greater # of electrons corresponds to greater polarizability, which makes for stronger
LDFs
Every molecule has LDF forces regardless of if they have other IMFs, be it dipole dipole or
H bonding.

Molecules with weak IMFs ( 4,5,6 above. ) tend to be gases at room temp (Generally have
ONLY LDFs)
Molecules with stronger intermolecular forces tend to be liquid at room temp. ( Strong
LDFs, DD, HB. )
Ionic substances do not have Intermolecular forces. ( Usually solid at room temp )
If a molecule in a liquid builds up enough energy it can break free into a vapor, this is
vaporization. ( This is also responsible for vapor pressure. )
Single < Double < Triple bonds in bond strength
Single > Double > Triple in bond length

Lewis Structures + VSEPR:

It is important to remember that lewis structures

are a visual representation of bonds and electrons.

For single atoms/ions all valence electrons should be drawn in. It is important to note that
if something is ionized and has no electrons left, then no dots are drawn in - do not go
down to the next shell! Also, ions should be put inside of brackets with the charge
indicated in the upper right corner.


In molecules, bonds are represented by drawing a line between the two atoms. These lines
can be used to represent single, double or triple bonds between two atoms in a molecule.

An important thing to remember is that each bond represents two electrons, any unbonded
electrons should be shown in pairs (two dots) ( with the exception of an odd total for

15
valence electrons ) and that the total number of electrons should be equal to that of the
sum of the valence electrons of the atoms. ( Make sure to add/subtract the charge from
this total if working with polyatomic ions! ) (Negative adds electrons, positive subtracts
electrons!)

Lastly, as with single atoms, molecules should be drawn inside of brackets if they are ions.


VSEPR ( Valence Shell Electron Pair Repulsion ) Theory : A model used to predict the
geometry ( structure ) of molecules based on bonds and lone pairs of electrons. The idea is
that electron pairs will arrange themselves in a way so as to minimize the repulsion
between them.

Electron Domain Geometry : The shape of the molecule that arises from the configuration
of the molecule about the central atom. The geometry is determined by the number of
electron domains ( Also called Steric Number. ) regardless of whether they are bonds or
lone pairs of electrons. Often referred to as hybridization (only sigma bonds hybridize).

Molecular Geometry : A more specific shape of the molecule that arises from the
configuration of the molecule about the central atom. The geometry is determined by the
number of bonds as opposed to the total number of electron domains.

***BOND ANGLES ARE DETERMINED BY THE ELECTRON DOMAIN GEOMETRY!

***



REACTIONS

Assume molecules without (solid, liquid, or gas) are aqueous (dissolve in water)

Synthesis- When two simple compounds form to make a more complex one
2H2 + O2 → 2H2O (liquid)

Decomposition- opposite of Synthesis, complex compound breaking into simple

compounds, usually requires energy (heat), usually endothermic
2MgO → 2Mg + O2

f-Base- when an acid(H+) reacts with a base(OH-) to form a salt and water

16
 HCl + NaOH → NaCl + H2O(liquid)

Oxidation-Reduction- reaction that changes the phase of a species ( Covered in greater

detail later on ) (involves transfer of electrons, used in batteries)
Cu+ + e- → Cu(solid)

Precipitation - Two aqueous solutions mix, sometimes a precipitate (solid excess

compound) is formed, also a cation and anion mixing can create a salt precipitate (non
aqueous)
K2CO3 + Mg(NO3)2 → 2(KNO3) + MgCO3(solid)
Also since the K+ and NO3- both start out aqueous and end aqueous, they don't
really do much and are called spectator ions

Net ionic equation - break up the molecules (ex. K2CO3 into 2K+ + CO32-), then remove
all spectator ions from the equation

The AP test will provide you with most solubility rules, the ones it will not tell you are:
Compounds with (positive) Alkali metals are always soluble (Li+, Na+, K+, Rb+, Cs+, Fr+)
Compounds with NO3- are always soluble.
Compounds with NH4 are always soluble.
Chloride ( As well as Bromide and Iodide ) compounds are usually soluble, the exceptions
being when paired with Ag+, Hg+, and Pb+2
Compounds with SO4 -2 are usually soluble, the exceptions being when paired with Ba+2,
Sr+2, Hg+2, and Pb+2
Hydroxides are insoluble except when in a compound with any of the first three bullets.

BALANCING CHEMICAL EQUATIONS

Reactant side (starting compounds go here)→ Product side (compounds after reaction go
here)

When balancing an equation, all elements or compounds must be equal ( Conservation of

mass)
For example, 2(H2O) → 10(H2) + 5(O2) is incorrect, because the reactant side only has 4
hydrogens, and 2 oxygens
However 2(H2O) → 2(H2) + O2

REDOX

For a reaction to be a REDOX species in the reaction must undergo a change in charge.
Reduction is when electrons are gained.
Reduction is Gain: RIG

17
Gaining Electrons Reducing: GER
Oxidation is when electrons are lost.
Oxidation is Loss: OIL
Losing Electrons Oxidation: LEO
Just remember OIL RIG/LEO goes GER
If an atom is by itself on one side of the reaction, and then on the other side of the reaction
it is in a compound that should be an instant clue to you that this is a REDOX reaction,
because that means that the charge changes
Some Important rules for determining charge to follow are:
An element by itself always has a charge of 0 ( Even diatomics! )
Oxygen will always have a charge of -2 ( Unless by itself or if it is in the peroxide molecule
O2 -2 in which case each oxygen will have a charge of -1. )
Hydrogen will have a charge of +1. ( Unless by itself, or in the case that it has a -1 charge -
don’t worry about remembering this, it won’t happen unless they tell you it will. )
Group 1 and 2 elements usually have a +1 and +2 charge, respectively.
Group 17 elements will usually take on a -1 charge.
While group 1, 2 and 17 elements usually only have one oxidation state, transition metals
can have several.
The overall charge of the compound should be zero unless it is an ion, in which case it
should then add up to the indicated charge of the ion. ( Unless it is a polyatomic you should
have memorized, in which case you should have also memorized the charge - it may not be
given to you. )
Ex :
Reaction : AB + C A + CB
Let’s assume ‘A’ and ‘C’ are group one elements and ‘BC’ is a group 7 element. On the
reactant side A has a charge of +1, B has a charge of -1, and C is neutral. We know this
because A is a group 1, B is a group 17 and C is by itself. On the product side, A now has a
charge of 0, B still has a charge of -1, and C now has a charge of +1. We know this because
A is by itself, C is a group 1 element, and B is a group 17 element.
This means that:
A has gone from a charge of +1 to 0. It has been reduced, or has gained
electrons/become more negative. (why can’t electrons be positive reeeeee)
B has not undergone a change in charge - no change means it is not
part of the redox! ( This is important in Net Ionic and Half Reactions, defined and
explained below. )
C has gone from a charge of 0 to +1. It has been oxidized, or has
lost electrons/become more positive.

Balancing REDOX:

Usually the types of REDOX reactions we should expect to deal with are not so
simple as the example given above. They generally will require balancing the equation as
well, and when this is the case, they usually also need to be broken down into half
reactions so as to maintain both a conservation of charge and mass.

18
Full Ionic Equations represent all species present in the reaction, whether they
‘participate’ or not. Weak electrolytes are NOT split up into ions. ( These are also good to
use when determining precipitates if you are shaky on solubility rules. )
Example :
HF(aq) + NaOH(aq)NaF(aq) + H2O(l)* will become
HF(aq) + Na+(aq) + OH-(aq)Na+(aq) + F-(aq) + H2O(l)
*Note that this does not represent REDOX, it is simply here to
demonstrate the concept of Full/Net Ionic equations until I or someone else can find/think
of a better one.
Net Ionic Equations represent all species that undergo a change during the
reaction, whether it be in charge, phase or if it was/becomes part of a weak electrolyte. As
with before, weak electrolytes remain unsplit. Anything that does not undergo a change is
omitted*. These are called spectator ions.
Example : HF(aq) + OH-(aq)F-(aq) + H2O(l)
*Note that Na+(aq) is not in the net ionic yet F-(aq) is despite
Fluorine being on both sides in aqueous form. This is because Na+(aq)
appears the same way on both sides, but fluorine is part of a weak acid (And thus, a weak
electrolyte ) HF(aq) as a reactant, but appears by itself as F-(aq) on the product side. It is
because it does not appear the exact same way twice that it is included.
By looking at the reaction and the Net Ionic it becomes easier to determine what the half
reactions will be since it is easier to see what is reduced and what is oxidized. ( A single
species can be both reduced AND oxidized, usually a reactant. )

ELECTROCHEMISTRY

https://www.youtube.com/watch?v=A0KaYHpPxqg
This is pretty helpful for Electrochem

2 Of Amerikz Most Wanted

Two major ‘cell’ types

Galvanic, or voltaic, thermodynamically favored

19
Electrolytic, thermodynamically UNfavored (has to use another power source, such as a
battery)
Depends on the ∆G
If ∆G is less than 0 (G<0) it is FAVORED
If ∆G is greater than 0 (G>0) UNFAVORED

Electrodes
Anode
Electrode that gets oxidized
Seems to get smaller

Cathode
Electrode that get reduced
Seems to get larger

Salt bridge
Used to maintain an electric neutrality in a galvanic cell
Anions from the salt bridge flow towards the anode
Cations from the salt bridge flow towards the cathode
Voltmeter is used to measure the cell potential
Electron always flow from anode to cathode
Mnemonic devices
AN OX (Anions Oxidize)
RED CAT (Reductions are for Cations)
FAT CAT (Electrons flow From the Anode To the CAThode)
EPA
Electrolysis has a Positive Anode since the polarities are reversed!
AN OX chases the RED CAT

Cell potential
Positive cell potential makes the cell work
E⁰ = E(reduction) - E(oxidation) (Hess’s Law type question: change the sign of oxidation
half-reaction because it’s reversed in the overall redox reaction)
⁰ in E⁰ means standard conditions: 298 K, 1 atm, 1.0M in each solution
E⁰ is intensive, so it doesn’t change with mass. Multiplying a half-reaction by 2 to balance
the redox reaction does NOT mean you multiply the E value.
If both reduction potentials positive: the lower potential will probably become the oxidation
reaction
If one positive, one negative: the negative one will probably become the oxidation reaction
If they are both negative, the lower negative one will flip and become the oxidation
reaction
E(cell) is E⁰ at different temperatures, pressures, and molarities

20
Nernst Equation: E(cell) = E⁰ - RT/nF (ln Q)
E⁰ is standard reduction potential
R = 8.314 J / mol K
n= moles of electrons transferred (balance redox to see this)
F = 96485 Coulombs per mol of e-
Q = reaction quotient, molarities of product over molarities of reactants (see
equilibrium)
Nernst Equation at 298 K: E(cell) = E⁰ - 0.0592/n ln(Q)
Nernst Equation at Equilibrium: E(cell) = -RTln(k)
At equilibrium, the standard cell stops producing voltage so E⁰ becomes 0

Concentration cell: a cell with the same elements on either side, driven by molarity
difference
Galvanic Cell Notation: Cu | Cu2+ || Cd2+ | Cd
Where Cu(s) is being oxidized in the anode and Cd(s) is being reduced in the cathode
Electromotive force to pull electrons as they move from anode to cathode
Known as E(cell)
Measured in Volts
1V (volt)= 1 Joule/ 1 coulomb
Electrolysis - dimensional analysis questions.
Amps are C/second
Faraday constant 96485 is C/mol e-
Sec * (C/second) * (mol e-/C) * (mol A/mol e-) * (g A/mol A)
Can be applied backward to go from grams or moles of A to time
Just fencepost it out

∆G, Gibbs Free Energy (kJ/mol)

∆G=∆H-T∆S
∆H=Enthalpy Change, T= Temperature, ∆S= Entropy Change
Important to remember that ∆H is normally in kJ and ∆S is normally in J, so you
need to convert it to J for the equation to work
∆G=-nFE
F=96,485 C per mole of e- , n=mol of electrons transferred (in balanced full equation,
usually harder than you’d expect to find), E=reduction potential (Ecell)

Stoichiometry
When dealing with trying to find out information on either the reactant or product side for
questions, follow these steps
Convert to moles (grams/molar mass)
Find the limiting reactant (which material runs out first (not always what is the lowest))
Use the balanced equation to see how many moles it will need

21
Convert moles into desired unit

ENTHALPY

Enthalpy- measure of how much energy is released or absorbed

Energy is basically heat, so it is good to relate it to heat

Basic Rules of enthalpy

When Bonds are MADE, energy is RELEASED
When Bonds are BROKEN, energy is ABSORBED
It takes energy to break bonds (think that atoms are attracted to each other
while in bonds, so you need to push them away from each other), making bonds releases
energy
Enthalpy is usually written as ΔH
When ΔH is negative (-ΔH) heat is being RELEASED and it is Exothermic
Heat is treated as product (for example a candle flame)
When ΔH is Positive (ΔH) Heat is being ABSORBED and it is endothermic
Heat treated as reactant (ex: melting an ice cube)
Exothermic- heat (energy, or electrons) is released (hot stuff) ( PE of Reactants should be
higher than PE of products! )
** More likely to be spontaneous
Endothermic- heat( energy, or electrons) is absorbed (cold stuff) ( PE of reactants should
be lower than PE of products! )

MEMORIZE THESE
Activation energy- is how much energy is needed to fully break all the bonds in the
reactants (none form in the products yet)
The top of the peak is the activated complex, and no bonds are present when at it,
immediately after product bonds form, immediately before, reactant bonds break

Catalyst- a substance that speeds up a reaction, while not affecting its outcomes, its more
of a short cut for the reaction (lowers it for both forward and reverse)
It speeds up a reaction by providing a lower activation energy without changing the
products or delta H.


KINETICS

RATE LAWS

Rate Laws- The generic equation for a rate law is Rate = k[A]^x[B]^y[C]^z, with x being
the exponent determined by the ‘order’ of [A], y for [B], and z for [C]. Rate laws give us the
rate/‘speed’ of the reaction in relevance to the concentrations of the reactants in the ‘rate

22
determining step.’

Rate determining/Slow step- This the part of the reaction that is responsible for how long it
will take overall - aka it determines the rate.

Fast Step(s) - Usually we can just ignore these since they are assumed to happen so
quickly that they have a negligible impact on the overall reaction time.

ZERO ORDER: FIRST ORDER: SECOND ORDER:

Zero Order Rates- [A] vs time graph linear


First Order Rates- ln[A] vs time graph linear; half life independent of concentration

Second Order Rates- 1/[A] vs time graph linear

Rate Constant k- increases with temperature; units for k found by substituting the units (M
for [A] and M/s for Rate) into the rate law than solving for k.

Half Life : t½ = 0.693/k (only for first order) ( In the reference table! ) OR can be
determined from a graph if given
Concentration vs Time - even if it is first order! ( Just look for where the concentration is
half of its original value, and to double check your answer, either plug it into the equation
and solve for k, or check from the half value and halve it again to see if it is the same
amount of time. )

- The half life of a reaction is the time it takes for the concentration of the
reactants/species in question to reach half of its original value. ( Ex : In a generic equation
A + B C where [A] is 1.00 M, the value for the half life would be how long it will take for
[A] to decrease from 1.00 M to 0.50 M, from 0.50 M to 0.25 M, and so on and so forth. )

- Half life remains constant throughout the reaction (Only if First-order)! It should not be
changing - if you check your answer at different intervals on a graph, and it has changed
by a significant amount, or if the half life equation and graph derived values are very
different you are doing it wrong! ( Sometimes finding an answer by the graph and finding
the answer using k can give slightly different values. The same goes for the reverse, if
using half life to solve for k. Don’t worry if the answer is only changed by .005 or so. )

Beer-Lambert Law- A=ɛbc

A - Absorbance

23
 -
ɛ- molar absorptivity (usually a constant) - specific to the solute and optimal
wavelength
b- path/cuvette length/width
c- concentration

Reaction Mechanisms

Collision theory
A postulate that states that a reaction occurs when the reactant molecules
collide with sufficient energy and proper orientation.
There are 5 ways to increase the speed at which the reaction occurs
1. Temperature increase: Means the average kinetic energy of the molecules is higher so
the molecules are more likely to collide with enough energy to overcome the activation
energy.
2. Surface area increase
3. Catalysts
4. Increasement of concentration: More reactant molecules present means more collisions
which in turn raises the likelihood of having successful collisions
5. Pressure increase: Less space for reactant molecules to move around so they hit each
other more frequently, resulting in more successful collisions overall

THERMODYNAMICS

ΔH= Enthalpy (Heat/energy)

ΔS= Entropy Change (how chaotic something is)
ΔG= How much free energy a reaction has to give

Basic rules (@ Standard Conditions= °)

All gases are 1 atm
All liquids are pure
All solids are pure
All solutions are at 1 Molarity
Temperature is usually room temp 25 Celsius, 298 Kelvin
Element in normal state has zero energy

ΔH = Enthalpy
-ΔH is exothermic, more stable
+ΔH is endothermic, less stable
Pure element should have 0 ΔH

24
Entropy (ΔS)
Entropy of a perfect, pure crystal @0K=0
Universe prefers chaos (2nd law of thermodynamics)
Gas is the most chaotic state, aqueous being 2nd most chaotic
Solid is the least chaotic state
Increase in Entropy
Increasing in the number of moles of gas
Ex: A(g)+B(l)2 C(g)
Moving from a more ordered state to a more chaotic (solid→liquid→gas)
Negative ΔS means that the system is becoming more orderly, positive ΔS means that the
system is becoming more disorderly

Hess’s Law- overall enthalpy change in a reaction is the sum of all the reactions in the
process

q=mcΔT

Used to find heat of process where

q: thermal energy
m: mass
c: Specific heat→ amount of heat required to raise the temperature of 1g of substance by
1°C
ΔT: change in temperature (Tfinal-Tinitial)

Gibbs Free Energy Equation: ΔG= ΔH- TΔS


Enthalpy in solutions
Bonds between ionic substances require energy, but can usually be done with energy made
through dipole attractions
This is explained through 3 steps, NaCl will be an example
Solvent Bonds Break, NaCl breaks into Na+ and Cl-,
H’s and O’s in H2O, which are normally attracted to each other (H+’s like to naturally be
closer to O-’s) will spread apart more THEY DO NOT BREAK
Then Na+ moves in between the molecules, and O-’s are attracted to it, so they huddle
around it, the H+’s huddle around Cl-

EQUILIBRIUM
Deals with reversible reactions (as opposed to going to completion)
A system is at equilibrium when the forward rate and the reverse rate are equal (products
and reactants continue to be produced)
The equilibrium constant Kc is the ratio of the product to reactant at equilibrium
Q tells the reaction progress
Q>K - reaction will proceed to the reverse direction
Q<K - reaction will proceed in the forward direction

25
Build timeline
Always shift towards equilibrium
____________K<------------Q_____________________________

LE CHATELIER’S PRINCIPLE
Whatever stress added to the system will be relieved by the system

***Changing pressure with inert gas will not shift the reaction
***there is an error in this chart - increasing temperature of system favors endothermic
reaction

Equilibrium and pressure

Known as Kp as is very interesting when applied to Le Chatelier's principle

Goes to the side with the fewest molecules, when P is increased**
Volume change is analogous to pressure change but they are inversely related
VERY IMPORTANT EQUATION!!!!
Kp = Kc(RT)∆n
Where R is the gas constant (0.08206 L*atm/mol*K) and deltaN is the difference in moles
between reactants and products. deltaN can be negative
This convert pressure to concentration and vice versa

ACIDS AND BASES

Strong acids ionize completely

HCl, HBr, HI, *H2SO4, HClO4, HNO3
*H2SO4 is polyprotic, only the first H ionizes completely, the second H is weak
All other acids are weak
Strong bases include hydroxides of Group 1A and most of Group 2A metals
Weak bases usually contain a nitrogenous base

Kw=1.0*10-14 at 25 OC
Ka * Kb = Kw
Kw is the water dissociation constant.
Since the ionization of water is an endothermic process, higher temperature increases Kw
and decreases the pH
Lower pH of pure water does not mean the the water becomes acidic at higher
temperature relates to the equation for the autoionization of water, 2H2O H3O + OH . As
temperature is raised, the pH of water decreases since it will begin to favor the products,
but it remains neutral because pOH also decreases due to the increasing concentration of
both ions. )
Neutrality is when [H3O+] = [OH-] ( NOT always 7! )

26
CONJUGATE ACIDS AND BASES

Bronsted-Lowry
The acid is the proton donor and the base is the proton acceptor
This proton is usually shown as an H+ ion, but since H+ does not exist simply by itself, it
piggybacks onto an H2O molecule to form the hydronium ion: H3O+ (H+ is acceptable in
AP)
Water is amphoteric (can act as either an acid or a base). Another name is amphiprotic
For weak acids, higher percent dissociation for lower concentrations
%dissociation= ([H+]/[HA]) x 100%
For diprotic and triprotic acids, each subsequent dissociation is less favored
(Ka1>Ka2>Ka3) Can be seen in H2SO4 where the first H+ leaves easily but the Ka for
HSO4- is much lower (~10^-2)
Oxyacids- more Oxygens will cause the electrons to be pulled toward oxygens and it is
easier for the H+ to ionize (more oxygens=stronger acid)
Lewis Model
Acids will accept an electron pair and bases will donate an electron pair.
Weak acids have a strong conjugate base and low Ka value.
Strong acids have weak conjugate base and have a high Ka value

BUFFER SOLUTIONS
Usually made of a weak acid and its conjugate base or vice versa
To make a buffer solution, choose a conjugate acid-base pair with pKa close to the target
pH
A buffer will counteract large changes in pH
Capacity of a buffer is dependent on the concentrations of the 2 components of the buffer
(higher the concentration, the higher the capacity)
pH= pKA+ log(Base/Acid) OR pOH=pKb + log [HB+]/[B]
k for the HH equ our teacher taught
[H+]=Ka(mol acid/mol base)
Ideal buffer is equal concentration acid/conj. Base and vice versa (pH=pKa)

TITRATION
If you have a solution with unknown concentration and a solution with KNOWN
concentration you can perform a titration. This involves putting the known solution (acid or
base) into a buret over a flask of your unknown solution.
For best results you want to use an acid-base indicator with a range close to the
equivalence point. This way you will be able to see the dramatic change in pH via the
change in the indicator color. The most common indicator is phenolphthalein which is
colorless in acid and pink in base.

Titrant: The substance with the known molarity which is used to find the molarity of the
other substance, usually added to the analyte (our teacher did show us a released problem
where analyte/substance w/unknown molarity was being added to the titrant)

27
Analyte: Substance which is being analyzed to find its molarity
Indicator: pH=pKa

TITRATION GRAPH
Equivalence point is where moles of both the acid and the base are equal in 1:1 ratio
Strong/strong combinations will always have an equivalence point of pH=7.
Weak/strong combinations will have equivalence points with pH higher or lower than 7,
leaning towards the stronger acid/base
In a weak acid vs strong base titration, at the half-equivalence point [HA]=[A-] and
pH=pKa (also known as a perfect buffer)
Also , if pKa<pH, then the deprotonated form will dominate...a question from last year’s
frq that many tripped on
(If the pH is lower than the pKa, then the compound will be protonated. If the pH is higher than
the pKa, then th e compound will be deprotonated)
pKa>pH = protonated
pKa<pH = deprotonated

pH of an acidic/basic salt after titration (at eq point)

Only strong acid and weak base (or vice versa) will form non-neutral salt
moles acid=moles base=moles of salt produced
If given Kb, find Ka and vice versa
Use RICE table to find concentration of salt at eq point
Use your common equilibrium formula to determine the pH

USEFUL RESOURCES
Course Exam and Description, with useful questions and information
2013 Updated AP Chemistry Compilation Exams
2013 Updated AP Chemistry Practice Exam
2014 AP Chemistry Practice Exam
2014 AP Chemistry Free Response, (Answers here)
2015 AP Chemistry Free Response. (Answers here)
2017 IPE Chemistry AP Exam

Important things to sneak into your calculator (which i will be doing):

nernst
vesper
Kp=Kc(RT)^∆n
order reactions
E = E° - (0.0592/n)log Q
arrhenius equation
Raoult's law
MM=D(RT)/P

28
increasing temp by 10 the rate doubles√
[NMI HDL] nevins mom is hella dumb lol (inside joke but it is for bonds strength)
a mnemonic device to remember the order of strength of bonds and IMFs -
network covalent, metallic, ionic, hydrogen bonding, dipole-dipole, london dispersion
Keeping these may help you but they may have no effect. Better safe than sorry!

( 6.022 x10^23 )

Convert Moles to Grams (vice versa)

Moles = (gram/molar mass)
Grams = (Moles x Molar Mass)

Moles of a gas are found with n=PV/RT ( Rearranged Ideal Gas Law. )
Most likely will occur at STP (Standard Temperature & Pressure)
Pressure = 1 atm
Temperature = 273 K (0 C)
At STP one mole of an ideal gas will be at 22.4 Liters.

P = Pressure in atm
V = Volume in litres
n = # of moles
❑
R = Gas Constant ( .0821 )
mol∗K
T = Temperature in Kelvin ( Take Celsius and add 273 to get Kelvin )(Test won’t give
Fahrenheit)

Moles = (molarity)(liters of solute)

Percent Composition- percent of mass of each element in a compound

Divide each element's molar mass by compounds molar mass
Ca(NO3)2 = 164.10
Calcium = 1 = 40.08 x 1 = 40.08
Nitrogen = 2 = 14.01 x 2 = 28.02
Oxygen = 6 = 16 x 6 = 96.00
Ca = 40.08/164.10 = 24.42%
N = 28.02/164.10 = 17.07%
O = 96/164.10 = 58.50%d

Empirical Formula

29
Empirical formula can be found by calculating the moles of each element present, and then
dividing all mole values by the smallest number found. The values given by the division will
be the empirical formula subscripts.

Electron Configurations

Nucleus is positive, pulling in electrons which are negative, closer they are stronger the
attraction is i.e More protons= more attraction, closer electrons= more attraction

Electrons are repelled by other electrons, an electron between a valence and nucleus
causes the valence to be weaker (called shielding)

Completed shells are stable,atoms strive for completed shells

Aufbau principle- when building electron configurations, electrons go into shells in order of
increasing energy

Pauli Exclusion Principle- two electrons in a shell will not spin the same way

Hund’s rule- when an electron is added it will take up the lowest energy orbital

Amount of energy from an electron is dependent on how far away it is from the nucleus

Bohr Model - Quantum model for the atom where electrons are arranged in specific orbits
around the atom - used this to predict the emission spectrum of hydrogen given the
electron distribution. Problem - good for predicting some emission spectrums but does not
work with most. Also, electrons do not “orbit” in a perfect, flat plane.

If an atom is exposed to electromagnetic levels greater than the ones exerted by the
electrons it holds the electrons can be ejected.

Electron Configuration w/ periodic table

30
Use Periods (horizontal) as first number, then letter of the group, then amount
Examples
Hydrogen: 1s^1
Helium: 1s^2
Lithium: 1s^2, 2s^1
Sodium: 1s^2, 2s^2 2p^6, 3s^1
You can also ‘shorthand’ by using the Noble gases as starting points
Argon: [Ne] 3s^2, 3p^6

Moving left to right from periodic table atomic radius (distance from nucleus to valence
electrons) decreases (Zeff increases while valence electrons are added to same energy level,
so overall the electrons are pulled in more)

Moving down a group, atomic radius increases (more shells mean the valence electrons are
further out)

Removing electron creates a positively charged cation (X1+), the energy to remove the first
electron from a gaseous ion is the ionization energy, the energy to remove the next
electron is the second ionization energy, which is always higher than the first ionization
energy as the number of protons remains the same while there is one less electron, so less
force repelling and the same force attracting, overall requiring more force to remove the
electron.

If the electron being removed drops a shell, ionization energy increases MASSIVELY.

Electronegativity

31
 Small atoms have larger electronegativity levels
Atoms close to having stable electron amounts have stronger electronegativity levels

BONDS

Bonds - sharing of electrons, more electrons shared = stronger bond

So single bonds (1 σ) have the least bond energy
Bond Energy Strength: triple (2 π and 1 σ)>double (1 π and 1 σ)>single (1 σ)

Ionic Bond- Bonds held by electrostatic forces, one will be + (cation(red)) while the other is
- (anion(black))
Ionic bonds are where the anion completely removes the electron from the cation

Metallic Bond- Metals in the bond are usually stationary and share a sea of electrons,
making them good conductors as they give up electrons easily,this also makes the metals
malleable

Metals can bond with each other and form alloys

Interstitial Alloy- metals with VERY different atomic radii combine
Example is carbon filling in the cracks around iron to make steel

Substitutional Alloy- Two metals with similar atomic radii combine

Example is Zinc and Copper, it is much more structured and uniform

32
Covalent Bonds- Two atoms share electrons and both count the electrons as part of their
valence shell (nonmetals)
The first covalent bond is called the sigma bond (σ), and has the lowest bond energy
and longest bond length, while the following bonds (2nd and 3rd) are called pi bonds

Network Covalent Bonds-these bonds are formed when a group of atoms share covalent
bonds into one big electron sharing cluster, it is very hard to break and has a high melting
point
Good contenders are Silicon and Carbon since they only have 4 valence electrons
(ex. SiO4)

Polarity- Atoms usually will have a electronic imbalance of sorts due to uneven distribution
of electrons (very based on symmetry)
Dipole- The most electronegative element will pull on electrons but not enough to become
ionic, but will keep them on one side of the atom, H 2O for instance has 2 H+ on one side,
connected to O- this makes the O- side negative, and the H+ side positive, more charges
means more dipole moment, which is how a dipole is measured
Dipole moments can attract each other, causing the H+’s in H2O to be attracted (they do
not form a bond) to the O-’s in other water molecules (this is a hydrogen bond described
below)

Intermolecular forces- forces that exist between 2 or more covalent molecules that hold
them together. (ex. London Dispersion forces, dipole/hydrogen bonds.) Stronger IMFs =
higher boiling point, lower vapor pressure

Hydrogen Bonds- Special type of dipole moment attraction that involves a positively
charged hydrogen to bond to a negative end of another molecule, usually one with high
electronegativity (H―F, H―O, H―N).
Extremely strong bond as hydrogen gives up its only electron so it uses all its strength to
hold onto it, molecules with these bonds have high boiling/melting points. A group of H-F
molecules, or any molecule w hydrogen bonding,are oriented so that the very positive H of
one molecule is attracted to the very electronegative F of another molecule.

London Dispersion forces- With large amounts of electrons moving in random patterns it is
possible for a majority to move to a specific side, creating a quick bit of polarity ( Dipole
Moment ), which can in turn disrupt other molecules. More electrons means more chance
of them congregating ( more polarizable ), so molecules with high boiling points are the
most vulnerable. Larger molecules are more polarizable so have more electrons and will
more likely have London dispersion forces
To compare the forces, compare the IMFs - larger is stronger

Bond order = (# of sigma and pi bonds/ # of atoms)

Formal Charge = Group Number - Lone electrons - # of sigma and pi bonds
(Group Number - dots - sticks = formal charge)

33
 - most stable structures generally have negative formal charges on the more
electronegative atoms and positive formal charges on the less electronegative
atoms

Bond Strength + Intermolecular Forces ( IMFs )

Strongest to weakest

[1.] Covalent Bonds

[a.] ( Usually liquids. )
[b.] Network Covalent Bonds are strongest possible (think diamond).
[c.] (sharing valence electrons means you need to break 2 sets of valence
electrons, making it the toughest to break)
[d.] (Melting/boiling point is lower than ionic)
[e.] Use 4,5,6 bonds below to base on other covalent bonds strength
[f.] Electrons are shared.
[g.] Nonmetal-Nonmetal.

[2.] Ionic Bonds

[a.] Usually solid.
[b.] Energy to break based on coulombic attraction
[c.] Bad conductors in solid form as electrons don’t move much.
[d.] Generally good conductors in aqueous solution.
[e.] Electrons are transferred.
[f.] Metal-Nonmetal.
[3.] Metallic Bonds
[a.] Transition metals seem to be strongest of all metals
[b.] High melting point, usually below ionic
[c.] Generally good conductors as a solid. ( ‘Mobile sea of electrons.’ )
[d.] Metal only.
[4.] Hydrogen Bonds ( Actually an IMF despite ‘bond’ being in the name. )
[a.] Special/specific kind of Dipole attraction
[b.] Most commonly between H and F,O,or N ( “Hydrogen bonding is FON!” )-
corresponding F,O, or N must have a pair of unpaired e-
[c.] Strongest IMF.
[d.] Contains both LDFs and DD.
[5.] Non-Hydrogen Dipole bonds ( Dipole-Dipole - DD)
[a.] Usually long lasting
[b.] Stronger than LDFs but weaker than HB.
[c.] Contains LDFs
[d.] Polar
[6.] London Dispersion Forces ( Sometimes called Induced Dipole )
[a.] Only creates a dipole temporarily
[b.] Weakest IMF.

34
 [c.] Is the reason why larger molecules tend to have higher boiling points/lower
vapour pressure than smaller ones, can be used to help determine which
molecules have the stronger IMFs if they both have D-D and/or HB.
[d.] Greater # of electrons corresponds to greater polarizability, which makes for
stronger LDFs
[e.] Every molecule has LDF forces regardless of if they have other IMFs, be it
dipole dipole or H bonding.

a
Molecules with weak IMFs ( 4,5,6 above. ) tend to be gases at room temp (Generally have
ONLY LDFs)
Molecules with stronger intermolecular forces tends to be liquid at room temp. ( Strong
LDFs, DD, HB. )
Ionic substances do not have Intermolecular forces. ( Usually solid at room temp )
If a molecule in a liquid builds up enough energy it can break free into a vapor, this is
vaporization. ( This is also responsible for vapor pressure. )
Single < Double < Triple bonds in bond strength
Single > Double > Triple in bond length

Lewis Structures + VSEPR:

It is important to remember that lewis structures are a visual representation of bonds and
electrons.

For single atoms/ions all valence electrons should be drawn in. It is important to note that
if something is ionized and has no electrons left, then no dots are drawn in - do not go
down to the next shell! Also, ions should be put inside of brackets with the charge
indicated in the upper right corner.

35
In molecules, bonds are represented by drawing a line between the two atoms. These lines
can be used to represent single, double or triple bonds between two atoms in a molecule.

An important thing to remember is that each bond represents two electrons, any unbonded
electrons should be shown in pairs (two dots) ( with the exception of an odd total for
valence electrons ) and that the total number of electrons should be equal to that of the
sum of the valence electrons of the atoms. ( Make sure to add/subtract the charge from
this total if working with polyatomic ions! ) (Negative adds electrons, positive subtracts
electrons!)

Lastly, as with single atoms, molecules should be drawn inside of brackets if they are ions.

VSEPR ( Valence Shell Electron Pair Repulsion ) Theory : A model used to predict the
geometry ( structure ) of molecules based on bonds and lone pairs of electrons. The idea is
that electron pairs will arrange themselves in a way so as to minimise the repulsion
between them.

Electron Domain Geometry : The shape of the molecule that arises from the configuration
of the molecule about the central atom. The geometry is determined by the number of
electron domains ( Also called Steric Number. ) regardless of whether they are bonds or
lone pairs of electrons. Often referred to as hybridization (only sigma bonds
hybridize).

Molecular Geometry : A more specific shape of the molecule that arises from the
configuration of the molecule about the central atom. The geometry is determined by the
number of bonds as opposed to the total number of electron domains.

***BOND ANGLES ARE DETERMINED BY THE ELECTRON DOMAIN GEOMETRY!

***

36
37
38
39
REACTIONS

Assume molecules without (solid, liquid, or gas) are aqueous (dissolve in water)

Synthesis- When two simple compounds form to make a more complex one
2H2 + O2 → 2H2O (liquid)

Decomposition- opposite of Synthesis, complex compound breaking into simple

compounds, usually requir as es energy (heat), usually endothermic
2MgO → 2Mg + O2

Acid-Base- when an acid(H+) reacts with a base(OH-) to form a salt and water
HCl + NaOH → NaCl + H2O(liquid)

Oxidation-Reduction- reaction that changes the phase of a species ( Covered in greater

detail later on ) (involves transfer of electrons, used in batteries)
Cu+ + e- → Cu(solid)

Precipitation - Two aqueous solutions mix, sometimes a precipitate (solid excess

compound) is formed, also a cation and anion mixing can create a salt precipitate (non
aqueous)
K2CO3 + Mg(NO3)2 → 2(KNO3) + MgCO3(solid)
Also since the K+ and NO3- both start out aqueous and end aqueous, they don't
really do much and are called spectator ions

Net ionic equation - break up the molecules (ex. K2CO3 into 2K+ + CO32-), then remove all
spectator ions from the equation

The AP test will provide you with most solubility rules, the ones it will not tell you are:
● Compounds with (positive) Alkali metals are always soluble (Li+, Na+,K+ , Rb+,
Cs+, Fr+)
● Compounds with NO3- are always soluble.
● Compounds with NH4 are always soluble.
● Chloride ( As well as Bromide and Iodide ) compounds are usually soluble, the
exceptions being when paired with Ag+, Hg+, and Pb+.
● Compounds with SO4 -2 are usually soluble, the exceptions being when paired with
Ba+2,
Sr+2, Hg+2, and Pb+2
● Hydroxides are insoluble except when in a compound with any of the first three
bullets.
● CAANGO (chlorates, acetates, ammonium, nitrates, group one)

BALANCING CHEMICAL EQUATIONS

40
Reactant side (starting compounds go here)→ Product side (compounds after reaction go
here)

When balancing an equation, all elements or compounds must be equal ( Conservation of

mass)
For example, 2(H2O) → 10(H2) + 5(O2) is incorrect, because the reactant side only
has 4 hydrogens, and 2 oxygens
However 2(H2O) → 2(H2) + O2

REDOX

● For a reaction to be a REDOX species in the reaction must undergo a change in

charge.
● Reduction is when electrons are gained.
○ Reduction is Gain: RIG
● Oxidation is when electrons are lost.
○ Oxidation is Loss: OIL
● Just remember OIL RIG
● If an atom is by itself on one side of the reaction, and then on the other side of the
reaction it is in a compound that should be an instant clue to you that this is a
REDOX reaction, because that means that the charge changes
● Some Important rules for determining charge to follow are:
○ An element by itself always has a charge of 0 ( Even diatomics! )
○ Oxygen will always have a charge of -2 ( Unless by itself or if it is in the
peroxide molecule O2 -2 in which case each oxygen will have a charge of -1. )
○ Hydrogen will have a charge of +1. ( Unless by itself, or in the case that it has
a -1 charge - don’t worry about remembering this, it won’t happen unless they
tell you it will. )
○ Group 1 and 2 elements usually have a +1 and +2 charge, respectively.
○ Group 17 elements will usually take on a -1 charge.
○ While group 1, 2 and 17 elements usually only have one oxidation state,
transition metals can have several.
○ The overall charge of the compound should be zero unless it is an ion, in
which case it should then add up to the indicated charge of the ion. ( Unless it
is a polyatomic you should have memorized, in which case you should have
also memorized the charge - it may not be given to you. )
○ Ex :
Reaction : AB + C A + CB
Let’s assume ‘A’ and ‘C’ are group one elements and ‘B’ is a group 17
element. On the reactant side A has a charge of +1, B has a charge of -1, and
C is neutral. We know this because A is a group 1, B is a group 17 and C is by
itself. On the product side, A now has a charge of 0, B still has a charge of -1,

41
 and C now has a charge of +1. We know this because A is by itself, C is a
group 1 element, and B is a group 17 element.
This means that:
A has gone from a charge of +1 to 0. It has been reduced, or has
gained electrons/become more negative. (why can’t electrons be positive reeeeee)
B has not undergone a change in charge - no change means it is not
part of the redox! ( This is important in Net Ionic and Half Reactions, defined and
explained below. )
C has gone from a charge of 0 to +1. It has been oxidized, or has
lost electrons/become more positive.

Balancing REDOX:

● Usually the types of REDOX reactions we should expect to deal with are not so
simple as the example given above. They generally will require balancing the
equation as well, and when this is the case, they usually also need to be broken down into
half reactions so as to maintain both a conservation of charge and mass.
● Full Ionic Equations represent all species present in the reaction, whether
they
‘participate’ or not. Weak electrolytes are NOT split up into ions. ( These are also good
to use when determining precipitates if you are shaky on solubility rules. )
○ Example :
HF(aq) + NaOH(aq) NaF(aq) + H2O(l)* will become
HF(aq) + Na+(aq) + OH-(aq) Na+(aq) + F-(aq) + H2O(l)
*Note that this does not represent REDOX, it is simply here to
demonstrate the concept of Full/Net Ionic equations until I or someone
else can find/think of a better one.
● Net Ionic Equations represent all species that undergo a change during the
reaction, whether it be in charge, phase or if it was/becomes part of a weak electrolyte. As
with before, weak electrolytes remain unsplit. Anything that does not undergo a
change is omitted*. These are called spectator ions.
○ Example : HF(aq) + OH-(aq) F-(aq) + H2O(l)
*Note that Na+(aq) is not in the net ionic yet F-(aq) is despite
Fluorine being on both sides in aqueous form. This is because Na+(aq)
appears the same way on both sides, but fluorine is part of a weak acid
(And thus, a weak electrolyte ) HF(aq) as a reactant, but appears by itself as
F-(aq) on the product side. It is because it does not appear the exact
same way twice that it is included.
● By looking at the reaction and the Net Ionic it becomes easier to determine
what the half reactions will be since it is easier to see what is reduced and
what is oxidized. ( A single species can be both reduced AND oxidized, usually
a reactant. )

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ELECTROCHEMISTRY

https://www.youtube.com/watch?v=A0KaYHpPxqg
This is pretty helpful for Electrochem

Two major ‘cell’ types

● Galvanic, or voltaic, thermodynamically favoured
● Electrolytic, thermodynamically UNfavoured (has to use another power source, such
as a battery)
○ Depends on the ∆G
■ If ∆G is less than 0 (G<0) it is FAVOURED
■ If ∆G is greater than 0 (G>0) UNFAVOURED
● Electrodes
○ Anode
■ Electrode that gets oxidized
● Seems to get smaller
○ Cathode
■ Electrode that get reduced
● Seems to get larger
● Salt bridge

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 ○ Used to maintain an electric neutrality in a galvanic cell
○ Anions from the salt bridge flow towards the anode
○ Cations from the salt bridge flow towards the cathode
● Voltmeter is used to measure the cell potential
● Electron always flow from anode to cathode
● Mnemonic devices
○ AN OX (Anions Oxidize)
○ RED CAT (Reductions are for Cations)
○ FAT CAT (Electrons flow From the Anode To the CAThode)
○ EPA
■ Electrolysis has a Positive Anode since the polarities are reversed!
○ AN OX chases the RED CAT

● Cell potential
○ Positive cell potential makes the cell work
○ E⁰ = E(reduction) - E(oxidation) (Hess’s Law type question: change the sign of
oxidation half-reaction because it’s reversed in the overall redox reaction)
■ ⁰ in E⁰ means standard conditions: 298 K, 1 atm, 1.0M in each solution
■ E⁰ is intensive, so it doesn’t change with mass. Multiplying a half-
reaction by 2 to balance the redox reaction does NOT mean you
multiply the E value.
■ If both reduction potentials positive: the lower potential will probably
become the oxidation reaction
■ If one positive, one negative: the negative one will probably become the
oxidation reaction
■ If they are both negative, the lower negative one will flip and become
the oxidation reaction
○ E(cell) is E⁰ at different temperatures, pressures, and molarities
■ Nernst Equation: E(cell) = E⁰ - RT/nF (ln Q)
● E⁰ is standard reduction potential
● R = 8.314 J / mol K
● n= moles of electrons transferred (balance redox to see this)
● F = 96485 Coulombs per mol of e-
● Q = reaction quotient, molarities of product over molarities of
reactants (see equilibrium)
■ Nernst Equation at 298 K: E(cell) = E⁰ - 0.0592/n ln(Q)
■ Nernst Equation at Equilibrium: E(cell) = -RTln(k)
● At equilibrium, the
● standard cell stops producing voltage so E⁰ becomes 0
○ Concentration cell: a cell with the same elements on either side, driven by
molarity difference
○ Galvanic Cell Notation: Cu | Cu2+ || Cd2+ | Cd
■ Where Cu(s) is being oxidized in the anode and Cd(s) is being reduced
in the cathode

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 ○ Electromotive force to pull electrons as they move from anode to cathode
■ Known as E(cell)
Measured in Volts
1V (volt)= 1 Joule/ 1 coulomb
● Electrolysis - dimensional analysis questions.
○ Amps are C/second
○ Faraday constant 96485 is C/mol e-
○ Sec * (C/second) * (mol e-/C) * (mol A/mol e-) * (g A/mol A)
○ Can be applied backward to go from grams or moles of A to time
○ Just fencepost it out
● ∆G, Gibbs Free Energy (kJ/mol)
○ ∆G=∆H-T∆S
■ ∆H=Enthalpy Change, T= Temperature, ∆S= Entropy Change
■ Important to remember that ∆H is normally in kJ and ∆S is
normally in J, so you need to convert it to J for the equation to
work
○ ∆G=-nFE
■ F=96,485 C per mole of e- , n=mol of electrons transferred (in balanced
full equation, usually harder than you’d expect to find), E=reduction
potential (Ecell)

When dealing with trying to find out information on either the reactant or product side for
questions, follow these steps
[1.] Convert to moles (grams/molar mass)
[2.] Find the limiting reactant (which material runs out first (not always what is
the lowest))
[3.] Use the balanced equation to see how many moles it will need
[4.] Convert moles into desired unit

ENTHALPY

Enthalpy- measure of how much energy is released or absorbed

Energy is basically heat, so it is good to relate it to heat

Basic Rules of enthalpy

When Bonds are MADE, energy is RELEASED
When Bonds are BROKEN, energy is ABSORBED
It takes energy to break bonds (think that atoms are attracted to each other
while in bonds, so you need to push them away from each other), making bonds releases
energy

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Enthalpy is usually written as ΔH
When ΔH is negative (-ΔH) heat is being RELEASED and it is Exothermic
- Heat is treated as product (for example a candle flame)
When ΔH is Positive (ΔH) Heat is being ABSORBED and it is endothermic
- Heat treated as reactant (ex: melting an ice cube)
Exothermic- heat (energy, or electrons) is released (hot stuff) ( PE of Reactants
should be lower than PE of products! )
** More likely to be spontaneous
Endothermic- heat( energy, or electrons) is absorbed (cold stuff) ( PE of Products

shou ld be lower than PE of reactants! )

MEMORIZE THESE
Activation energy- is how much energy is needed to fully break all the bonds in the
reactants (none form in the products yet)
The top of the peak is the activated complex, and no bonds are present when
at it, immediately after product bonds form, immediately before, reactant
bonds break

Catalyst- a substance that speeds up a reaction, while not affecting its outcomes,
it’s more of a short cut for the reaction (lowers it for both forward and reverse)
It speeds up a reaction by providing a lower activation energy without changing the
products or delta H.

KINETICS

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RATE LAWS

Rate Laws- The generic equation for a rate law is Rate = k[A]^x[B]^y[C]^z, with x
being the exponent determined by the ‘order’ of [A], y for [B], and z for [C]. Rate
laws give us the rate/‘speed’ of the reaction in relevance to the concentrations of
the reactants in the ‘rate determining step.’

Rate determining/Slow step- This the part of the reaction that is responsible for
how long it will take overall - aka it determines the rate.

Fast Step(s) - Usually we can just ignore these since they are assumed to happen so
quickly that they have a negligible impact on the overall reaction time.

ZERO ORDER: FIRST ORDER: SECOND ORDER:

Zero Order Rates- [A] vs time graph linear

First Order Rates- ln[A] vs time graph linear; half life independent of concentration

Second Order Rates- 1/[A] vs time graph linear

Rate Constant k- increases with temperature; units for k found by substituting the
units (M for [A] and M/s for Rate) into the rate law than solving for k.

❑
Half Life : t½ = k ( In the reference table! ) OR can be determined from a graph if
given Concentration vs Time - even if it is first order! (Just look for where the

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concentration is half of its original value, and to double check your answer, either
plug it into the equation and solve for k, or check from the half value and halve it
again to see if it is the same amount of time. )
- The half life of a reaction is the time it takes for the concentration of the
reactants/species in question to reach half of its original value. ( Ex : In a
generic equation A + B C where [A] is 1.00 M, the value for the half life would be
how long it will take for [A] to decrease from 1.00 M to 0.50 M, from 0.50 M to
0.25 M, and so on and so forth. )
- Half life remains constant throughout the reaction (Only if First-order)! It
should not be changing - if you check your answer at different intervals on a graph,
and it has changed by a significant amount, or if the half life equation and graph
derived values are very different you are doing it wrong! ( Sometimes finding an
answer by the graph and finding the answer using k can give slightly different
values. The same goes for the reverse, if using half life to solve for k. Don’t worry if
the answer is only changed by .005 or so. )

Beer-Lambert Law- A=ɛbc

A - Absorbance
-
ɛ- molar absorptivity (usually a constant) - specific to the solute and optimal
wavelength
b- path/cuvette length/width
c- concentration

Reaction Mechanisms

Collision theory
A postulate that states that a reaction occurs when the reactant
molecules collide with sufficient energy and proper orientation.
There are 5 ways to increase the speed at which the reaction occurs
● Temperature increase: Means the average kinetic energy of
the molecules is higher so the molecules are more likely to
collide with enough energy to overcome the activation
energy.
● Surface area increase
● Catalysts
● Increasement of concentration: More reactant molecules
present means more collisions which in turn raises the
likelihood of having successful collisions

48
 ● Pressure increase: Less space for reactant molecules to
move around so they hit each other more frequently,
resulting in more successful collisions overall

THERMODYNAMICS

ΔH= Enthalpy (Heat/energy)

ΔS= Entropy Change (how chaotic something is)
ΔG= How much free energy a reaction has to give

Basic rules (@ Standard Conditions= °)

● All gases are 1 atm
● All liquids are pure
● All solids are pure
● All solutions are at 1 Molarity
● Temperature is usually room temp 25 Celsius, 298 Kelvin
● Element in normal state has zero energy

ΔH = Enthalpy
-ΔH is exothermic, more stable
+ΔH is endothermic, less stable
Pure element should have 0 ΔH

Entropy (ΔS)
● Entropy of a perfect, pure crystal @0K=0
● Universe prefers chaos (2nd law of thermodynamics)
● Gas is the most chaotic state, aqueous being 2nd most chaotic
● Solid is the least chaotic state
● Increase in Entropy
○ Increasing in the number of moles of gas
■ Ex: A(g)+B(l) 2 C(g)
○ Moving from a more ordered state to a more chaotic (solid→liquid→gas)
● Negative ΔS means that the system is becoming more orderly, positive ΔS
means that the system is becoming more disorderly

Hess’s Law- overall enthalpy change in a reaction is the sum of all the reactions in
the process

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q=mcΔT
- Used to find heat of process where
- q: thermal energy
- m: mass
- c: Specific heat→ amount of heat required to raise the temperature of
1g of substance by 1°C
- ΔT: change in temperature (Tfinal-Tinitial)

Gibbs Free Energy Equation: ΔG= ΔH- TΔS

Enthalpy in solutions
Bonds between ionic substances require energy, but can usually be done with
energy made through dipole attractions
This is explained through 3 steps, NaCl will be an example
[1.] Solvent Bonds Break, NaCl breaks into Na+ and Cl-,
[2.] H’s and O’s in H2O, which are normally attracted to each other
(H+’s like to naturally be closer to O-’s) will spread apart more
THEY DO NOT BREAK

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 [3.] Then Na+ moves in between the molecules, and O-’s are
attracted to it, so they huddle around it, the H+’s huddle around
Cl-

EQUILIBRIUM
Deals with reversible reactions (as opposed to going to completion)
A system is at equilibrium when the forward rate and the reverse rate are equal
(products and reactants continue to be produced)
● The equilibrium constant Kc is the ratio of the product to reactant at
equilibrium
● Q tells the reaction progress
○ Q>K - reaction will proceed to the reverse direction
○ Q<K - reaction will proceed in the forward direction
○ Build timeline
Always shift towards equilibrium
____________K<------------Q_____________________________

[I.] LE CHATELIER PRINCIPLE

Whatever stress added to the system will be relieved by the system

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***Changing pressure with inert gas will not shift the reaction
***there is an error in this chart - increasing temperature of system favors
endothermic reaction

Equilibrium and pressure

● Known as Kp as is very interesting when applied to Le Chatelier's

principle
○ Goes to the side with the fewest molecules, when P is increased**
○ Volume change is analogous to pressure change but they are
inversely related
VERY IMPORTANT EQUATION!!!!

Kp = Kc(RT)∆n
Where R is the gas constant (0.08206 L*atm/mol*K) and deltaN is the
difference in moles between reactants and products. deltaN can be negative
This convert pressure to concentration and vice versa

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ACIDS AND BASES

Strong acids ionize completely

- HCl, HBr, HI, *H2SO4, HClO4, HNO3
- *H2SO4 is polyprotic, only the first H ionizes completely, the second H
is weak
- All other acids are weak
Strong bases include hydroxides of Group 1A and most of Group 2A metals

- Weak bases usually contain a nitrogenous base

Kw=1.0*10-14 at 25 OC
- Ka * Kb = Kw
- Kw is the water dissociation constant.
Since the ionization of water is an endothermic process, higher temperature
increases Kw and decreases the pH
- Lower pH of pure water does not mean the the water becomes acidic at
higher temperature relates to the equation for the autoionization of water,
2H2O H3O + OH . As temperature is raised, the pH of water decreases since
it will begin to favor the products, but it remains neutral because pOH also
decreases due to the increasing concentration of both ions. )
- Neutrality is when [H3O+] = [OH-] ( NOT always 7! )

CONJUGATE ACIDS AND BASES

Bronsted-Lowry
- The acid is the proton donor and the base is the proton acceptor
- This proton is usually shown as an H+ ion, but since H+ does not exist simply
by itself, it piggybacks onto an H2O molecule to form the hydronium ion: H3O+
(H+ is acceptable in AP)
- Water is amphoteric (can act as either an acid or a base). Another name is
amphiprotic
- For weak acids, higher percent dissociation for lower concentrations
- %dissociation= ([H+]/[HA]) x 100%
- For diprotic and triprotic acids, each subsequent dissociation is less favored
(Ka1>Ka2>Ka3) Can be seen in H2SO4 where the first H+ leaves easily but
the Ka for HSO4- is much lower (~10^-2)
- Oxyacids- more Oxygens will cause the electrons to be pulled toward oxygens
and it is easier for the H+ to ionize (more oxygens=stronger acid)
Lewis Model

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 - Acids will accept an electron pair and bases will donate an electron pair.
- Weak acids have a strong conjugate base and low Ka value.
- Strong acids have weak conjugate base and have a high Ka value

BUFFER SOLUTIONS
- Usually made of a weak acid and its conjugate base or vice versa
- To make a buffer solution, choose a conjugate acid-base pair with pKa close
to the target pH
- A buffer will counteract large changes in pH
- Capacity of a buffer is dependent on the concentrations of the 2 components
of the buffer (higher the concentration, the higher the capacity)
[]
- pH= pKA+ log( ) OR pOH=pKb + log [HB+]/[B]
[]
- k for the HH equ our teacher taught
- [H+]=Ka(mol acid/mol base)
- Ideal buffer is equal concentration acid/conj. Base and vice versa (pH=pKa)
TITRATION
- If you have a solution with unknown concentration and a solution with
KNOWN concentration you can perform a titration. This involves putting the
known solution (acid or base) into a buret over a flask of your unknown
solution.
- For best results you want to use an acid-base indicator with a range close to
the equivalence point. This way you will be able to see the dramatic change
in pH via the change in the indicator colour. The most common indicator is
phenolphthalein which is colorless in acid and pink in base.

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 - Titrant: The substance with the known molarity which is used to find the
molarity of the other substance, usually added to the analyte (our teacher did
show us a released problem where analyte/substance w/unknown molarity
was being added to the titrant)
- Analyte: Substance which is being analyzed to find its molarity
- Indicator: pH=pKa
- Useful Titration Calculator: Used to find the pH of any mixture, the volume or
concentration to reach a specific pH/equiv. point/midpoint. Also shows the
titration curve graph, and an optional step-by-step solution.

TITRATION GRAPH
- Equivalence point is where moles of both the acid and the base are equal in
1:1 ratio
- Strong/strong combinations will always have an equivalence point of pH=7.
- Weak/strong combinations will have equivalence points with pH higher or
lower than 7, leaning towards the stronger acid/base
- In a weak acid vs strong base titration, at the half-equivalence point
[HA]=[A-] and pH=pKa (also known as a perfect buffer)
- Also , if pKa<pH (> half equivalence point), then the deprotonated form will
dominate...a question from last year’s frq that many tripped on

pH of an acidic/basic salt after titration (at eq point)

- Only strong acid and weak base (or vice versa) will form non-neutral salt
- moles acid+moles base=moles of salt produced
- If given Kb, find Ka and vice versa
- Use RICE table to find concentration of salt at eq point
- Use your common equilibrium formula to determine the pH

DELICIOUSUSEFUL RESOURCES
● Course Exam and Description, with useful questions and information
● 2013 Updated AP Chemistry Compilation Exams
● 2013 Updated AP Chemistry Practice Exam
● 2014 AP Chemistry Practice Exam
● 2014 AP Chemistry Free Response, (Answers here)
● 2015 AP Chemistry Free Response. (Answers here)
● https://drive.google.com/drive/folders/1qNkNQfORqCUPrDDKsxfYKZx3DqbUIr5L
(This one ^ has all AP IPE tests from 2017, you’ll have to find the Chem one)
● https://www.adriandingleschemistrypages.com/ap/2017-ap-chemistry-exam-
draft-answers-and-comments/
● Detailed review over every unit
○ ^ includes videos, slides, notes, and quizzes over every topic/unit

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APPENDIX

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 ● Important things to sneak into your calculator (which i will be doing):

● nernst
● vesper
● Kp=Kc(RT)^∆n
● order reactions
● E = E° - (0.0592/n)log Q
● arrhenius equation
● Raoult's law
● MM=D(RT)/P
● increasing temp by 10 the rate doubles√
● [NMI HDL] nevins mom is hella dumb lol (inside joke but it is for bonds strength)
○ a mnemonic device to remember the order of strength of bonds and IMFs - network covalent,
metallic, ionic, hydrogen bonding, dipole-dipole, london dispersion
Keeping these may help you but they may have no effect. Better safe than sorry!
Going to fail this year bro.

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Dez nutz

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