AP Biology Review: Atomic Structure to Thermodynamics

AP Biology Review

I. Atomic Structure and Properties
II. Molecular and Ionic Compound Structure and Properties
III. Intermolecular Forces and Properties
IV. Chemical Reactions
V. Kinetics
VI. Thermodynamics
VII. Equilibrium
VIII. Acids and Bases
IX. Applications of Thermodynamics
X. Laboratory Overview

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I. Atomic Structure and Properties

Periodic table
Alkali metals, alkaline earth metals, transition metals, halogens, noble gases
Mass number = P + N
Isotopes - atoms of an element with different numbers of neutrons
Average atomic mass from weighted average of isotope mass and relative abundance (frequency)
Moles
PV = nRT
Avogadro’s number 6.022*10^23
AT STP (1 atm, 273K), 22.4 L/mol
Molarity M = moles/L
Percent composition - divide the mass of each element/compound by the total molar mass of the substance
Empirical formula is simplest ratio, molecular formula is actual formula for substance
Energy
Electron potential energy increases with distance from nucleus
Electron energy is quantized - can only exist at specific energy levels at specific intervals, not in between
Coulomb’s law: F = kq1q2/(r^2) where F is electrostatic force
Atoms absorb energy in the form of electromagnetic radiation as electrons jump to higher energy levels; when electrons drop levels (closer), atoms give off energy
Photoelectron spectroscopy
energy measured in electronvolts (eV)
Incoming radiation energy = binding energy + kinetic energy of the ejected electron
Electrons that are further away from nucleus require less energy to eject, thus will move faster
Photoelectron spectrum
Each section of peaks represents a different energy level (1, 2, 3, etc.)
Subshells within each energy level (shape of space electron can be found in orbiting nucleus) are represented by the peaks (1s, 2s, 2p, etc.)
s(2) - first subshell, p(6) - second subshell
Height of peaks determines number of electrons in subshell (ex. Peak of p subshell in energy level 2 will be 3x that of s subshell)
Electron configuration
Electron configuration - spdf - shorthand with noble gas first
Configuration rules
Aufbau principle - electrons fill lowest energy subshells available first
Pauli exclusion principle - 2 electrons in same orbital cannot have same spin
Hund’s rule - Electrons occupy empty subshells first
Zn +2, Ag +1, Al +3, Cd +2, most other transition metal charges vary
Periodic trends
Electrons are more attracted if they are closer to the nucleus, or if there are more protons
Electrons are repelled by other electrons - if there are electrons b/w the valence electrons and nucleus, the e- will be less attracted (shielding)
Completed shells are very stable, completed subshells are also stable; atoms will add/subtract valence electrons to complete their shell
INCREASING: atomic radius down left; ionization energy up right; electronegativity up right
Ionization energy - energy required to remove an electron from an atom
Electronegativity - how strongly the nucleus of an atom attracts electrons of other atoms in a bond
Electron affinity - energy change that occurs when an electron is added to an atom in the gas state (usually exothermic - energy is released)

II. Molecular and Ionic Compound Structure and Properties
Bonds
Atoms are more stable with full valence shells
Ionic bonds
Cation gives up electrons completely
Electrostatic attractions in a lattice structure
Metals and nonmetals (salts)
Coulomb’s law - greater charge leads to a greater bond/lattice energy (higher melting point)
If both have equal charges, smaller radius will have greater coulombic attraction
Ionic solid - electrons do not move around lattice; ionic solids are poor conductors of electricity; ionic liquids conduct electricity because ions are free to move around, though e- are still localized around particular atoms
Metallic bonds
Sea of electrons model - positively charged core is stationary while valence electrons are very mobile
Metals bond to form alloys - interstitial alloy w/ metals of different radii; substitutional alloy w/ metals of similar radii
Molecular covalent bonds
2 atoms share electrons - both atoms achieve complete outer shells
2 nonmetals
Creates molecules - 2+ atoms covalently bonded together
Single has 1 sigma bond - order 1, longest length, least energy; double has 1 sigma and 1 pi bond - order 2, int. length, int. energy; triple has 1 sigma and 2 pi bonds - order 3, shortest length, greatest bond energy
Bond forms when potential energy is at minimal level
Too close - potential energy is too high due to repulsive forces
Too far - potential energy is near 0 because attractive forces are very weak
Minimul PE occurs when repulsive and attractive forces are balanced
Network covalent bonds - lattice of covalent bonds - poor conductors, high melting and boiling points
Conductivity
Conductivity of different substances in different phases
Solid Aqueous Liquid Gas
Ionic No Yes Yes No
Molecular Covalent No No No No
Network Covalent No N/A No No
Metallic Yes N/A Yes No
Lewis dot structures
Resonance - for bond order calculations, average together all possible orders of a specific bond
BORON (B) is stable with 6 electrons - only one that does not need a full octet
Expanded octets - any atom of an element from n=3 or greater (those with a d subshell) can have [8,12] valence electrons on center atom
Noble gases form bonds by filling empty d orbital with electrons
Formal charge - number of valence electrons minus assigned electrons (1 e- for each line “shared” bond) - 0 for neutral molecules
Molecular geometry (VSEPR)
Double and triple bonds have more repulsive strength than single bonds - occupy more space
Lone electron pairs have more repulsive strength than bonding pairs, so molecules with lone pairs will have slightly reduced angles between terminal atoms
Hybridization - how many atoms are attached (sp, sp2, sp3, sp3d, etc.)
III. Intermolecular Forces and Properties
Polarity
Covalent bond where electrons are unequally shared - polar covalent
Dipoles are caused by polar covalent bonds - pair of opposite electric charges separated by some distance, like partial charges on atoms in a polar covalent bond
If 2 identical atoms bond (ex. Cl-Cl) the electrons are equally shared, creating a nonpolar covalent bond with no dipole
Bonds can be polar; so can molecules depending on the molecular geometry (and polarity of bonds - secondary)
In polar molecules, more electronegative atoms will gain negative partial charge; Usually central atom will be positive - exception is hydrogen (terminal), which is usually positive since it has less electronegativity
Intermolecular forces
Forces b/w molecules in a covalently bonded substance - need to be broken apart for covalent substances to change phases
Changing phase: ionic substances break bonds b/w individual ions; covalent substances keep bonds inside a molecule in place but break bonds b/w molecules
Dipole-dipole forces
Polar molecules - positive end of one molecule is attracted to negative end of another molecule
Greater polarity -> greater dipole dipole attraction -> larger dipole moment -> higher melting/boiling points
Relatively weak overall - melt and boil at low temps
Hydrogen bonds
Special type of dipole-dipole attraction where positively charged hydrogen end of a molecule is attracted to negatively charged end of another molecule containing an extremely electronegative element (F, O, N)
Much stronger than normal dipole-dipole forces since a hydrogen atom sharing/giving up its lone e- to a bond is left w/ no shielding
Higher melting/boiling points than substances held together only by other types of IMF
London dispersion forces
All molecules - very weak attractions due to random motion of electrons on atoms within molecules (instantaneous polarity)
Molecules w/ more e- experience greater LDF (more random motion)
Higher molar mass usually means greater LDF (as mass increases, e- increases for the molecule to remain electrically neutral)
IMF strength
Ionic substances are generally solids at room temp - melting them requires lattice bonds to be broken - energy determined by Coulombic attraction
Covalent substances (liquids at room temp) boil when IMF are broken; for molecules of similar size, from strongest to weakest: hydrogen bonds, permanent dipoles, LDF (temporary dipoles - greater for larger molecules)
Melting/boiling points of covalent substances are LOWER than for ionic substances
Bonding/Phases
Substances w/ weak IMF (LDF) tend to be gases at room temp (N2); substances w/ strong IMF (hydrogen bonds) tend to be liquids at room temp (H2O)
Ionic substances do not experience IMF - since ionic bonds are stronger than IMF, ionic substances are usually solids at room temp
Vapor pressure
Molecules in a liquid are in constant motion - if they hit the surface of the liquid with enough kinetic energy, they can escape the IMF holding them to other molecules and transition into the gas phase
Vaporization (NOT boiling) - no outside energy is added
Temperature and vapor pressure are directly proportional
At the same temp, vapor pressure is dependent on strength of IMF (stronger IMF, lower vapor pressure)
Solution separation
Solutes and solvents - like dissolves like
Paper chromatography
Piece of filter paper with substance on the bottom is dipped in water
More polar components of substance travel further up the filter paper with the polar water
Distance substance travels up the paper measured by retention/retardation factor Rf = (distance traveled by solute - substance being separated)/(distance traveled by solvent front - water)
Stronger attraction - larger Rf
Column chromatography
Column is packed with a stationary substance
separable solution (analyte) is injected, adhering to stationary phase
another solution (eluent - liquid/gas) is injected into column
more attracted analyte molecules will move through faster and leave column first
Distillation
Takes advantage of different boiling points of substances by boiling a mixture at an intermediate point
Kinetic molecular theory
Kinetic energy of a single gas molecule: KE = ½ mv^2
Average kinetic energy of a gas depends on the temperature (directly proportional), not the identity of the gas
Ideal gases have insignificant volume of molecules, no forces of attraction b/w molecules, and are in constant motion without losing KE
Deviations occur at low temperatures or high pressures
Volume becomes significant
Real gases have attractions and lower pressure than predicted
Maxwell-boltzman diagrams
Higher temp -> greater KE -> greater range of velocity
Smaller masses, greater velocities to have same KE
Effusion
Rate at which a gas escapes from a container through microscopic holes
High to low pressure
Greater speed, greater temp, greater rate of effusion
If at same temp, gas w/ lower molar mass will effuse first
Equations
Ideal gas equation: PV = nRT
R = 0.0821
Combined gas law: P1V1/T1 = P2V2/T2
Dalton’s law: P(total) = Pa + Pb + Pc + …
Partial pressure: Pa = P(total)*(moles of gas A)/(total moles of gas)
Density: D = m/V
From ideal gas law: Molar mass = DRT/P
Electromagnetic spectrum
E = hv; E = energy change; h = Planck’s constant 6.626*10^-34; v = frequency
C = lambda v; C = speed of light 2.99810^8; v = frequency; lambda = wavelength
Beer’s law: A = abc
A = absorbance; a = molar absorptivity; b = path length; c = concentration
Colorimetry - direct relationship between concentration and absorbance
IV. Chemical Reactions
Types of reactions
Synthesis: everything combines to form one compound
Decomposition: one compound + heat is split into multiple elements/compounds
Acid-base rxn: Acid + base -> water + salt
Oxidation-reduction (redox) rxn: changes the oxidation state of some species
Combustion: substance w/ H and C + O2 -> CO2 + H2O
Precipitation: aqueous solutions -> insoluble salt (+ more aq sometimes)
Can be written as net ionic - Those free ions not in net ionic are spectator ions
Solubility rules
Alkali metal cations or ammonium (NH4+) cations are ALWAYS soluble
Compounds with a nitrate (NO3-) anion are ALWAYS soluble
Common polyatomic ions
Calculations
Percent error: 100 * abs(experimental - expected)/(expected)
Combustion analysis - use law of conservation of mass (if x g of CO2 is produced, find g of C which will be starting amt)
Gravimetric analysis - when asked to determine the identity of a certain compound, find g of component produced, then use mass percent (g found / total sample mass) and compare to mass percent of options (molar mass of component / molar mass of entire compound)
Oxidation states
Neutral atoms not bonded to other atoms have an oxidation state of 0
Monoatomic ions have an oxidation state equal to the charge on that ion (ex. Zn2+ will be +2)
Oxygen is -2 (EXCEPTION: in hydrogen peroxide, H2O2, O is -1)
Hydrogen is +1 w/ nonmetals, -1 with metals
In absence of oxygen, most electronegative element in a compound will take an oxidation state equal to its usual charge (ex. F is -1 in CF4)
IF none of the above rules apply, determine the oxidation state by adding up all elements’ oxidation states to 0/charge on ion
C, N, S, P frequently vary oxidation states (low electronegativity)
Redox reactions
Write full rxn as 2 half reactions (oxidation and reduction; OIL RIG)
Add H2O to compensate for oxygen on one side
Add H+ to compensate for H from H2O on other side
Balance 2 half rxns to have the same number of electrons and add them together to produce one complete reaction
ACIDIC: stop here
BASIC: Add OH- to both sides - enough for all H+ on one side to be converted to H2O; then cancel out H2O so it only remains on one side
Acids and bases (briefly)
Color change signals the end of a titration (can be redox or acid/base)
Acids are capable of donating protons (H+); bases are capable of donating electrons
Water can act as an acid or base - amphoteric
V. Kinetics
Rate law
Rate = k [A]^x [B]^y [C]^z
Can calculate x, y, z via a table from (concentration factor)^x = (rate factor)
K is only dependent on temperature (always increases w/ T)
Keq = K1 (rate constant of forward rxn) / K2 (rate constant of reverse rxn)
K calculated by dividing any rate in table by the concentrations to their respective powers
Units for rate are M/s, units for conc are M -> calculate units for k from there
If A + 2B + C -> D; rate of formation of D = rate of disappearance of A and C = 0.5*rate of disappearance of B
Orders
Zero-order: Rate = k ; Concentration vs time has slope -k
First-order: Rate = k[A]; ln[A] vs time has slope -k; ln[A]t = -kt + ln[A]0
Second-order: Rate = k[A]^2; 1/[A] vs time has slope k; 1/[A]t = kt + 1/[A]0
Half-life: First order reactions only have a constant half life; t1/2 = ln(2)/k = 0.693/k
Collision theory
Chemical reactions occur because reactants are constantly moving and colliding with one another
When reactants collide with sufficient energy (activation energy Ea), a reaction occurs
Gaseous/aqueous: increased concentration increases rate of reaction
Stirring increases reaction rate for heterogeneous mixtures
Greater temp increases reaction rate
Reactions only occur if reactants collide with correct orientation
Reaction energy profile
Reaction mechanisms
Intermediates exist as species produced in a mechanism and consumed later
Adding up steps and canceling gives the overall equation
Elementary steps w/ 2 reactants are bimolecular; w/ 1 reactant are unimolecular
Speed is determined by slow step (rate-determining step)
Consistency is determined by slow step and those leading up to it
Slow step has highest activation energy
Catalysts
Catalysts increase rate without being consumed; do not appear in balanced equation
In a mechanism, catalysts enter first, then exit
Types: surface catalysis, enzyme catalysis, acid-base catalysis
VI. Thermodynamics
Temperature/heat
Temperature is the average amount of kinetic energy
Heat is the energy flow between two substances at different temperatures
First law of thermodynamics: energy can be neither created nor destroyed
When bonds form, energy is released; when bonds are broken, energy is absorbed
Exothermic - energy transferred from system to surroundings (delta H negative)
Endothermic - energy transferred from surroundings into system (delta H positive)
Energy diagrams
Enthalpy
Enthalpy of formation
Change in energy when one mole of a compound is formed from its component pure elements under standard conditions (25C/298K)
Delta Hf = delta Hf for products - delta Hf for reactants
If delta Hf is negative, energy is released when the compound is formed, so the product is more stable (exothermic)
If delta Hf is positive, energy is absorbed when the compound is formed, so the product is less stable than its constituent elements (endothermic)
Heat of formation is 0 when the pure element is in its standard state (ex. H2(g) or F2(g))
Bond energy
Delta H (J) = bond energies of reactants - bond energies of products
Multiply bond energies for each bond by the coefficient
Hess’s law
Finding delta H for the overall reaction from knowing delta H for the steps of the reaction
Flipping the equation flips the sign of delta H
Multiplying/dividing the equation by a coefficient multiplies/divides delta H by that coefficient
Adding/subtracting equations adds/subtracts their delta H values
Enthalpy of solution
Ionic substances dissolving in water
1: Breaking of solute bonds - energy required is equal to the lattice energy (positive delta H)
2: Separation of solvent molecules - water molecules must spread out to make room for the solute ions (positive delta H)
3: Creation new attractions - free floating ions are attracted to the dipoles of water molecules (negative delta H)
Hydration energy = step 2 + step 3 energies
Coulombic energy - increases with charge magnitude, decreases as size increases
Enthalpy of solution = step 1 + 2 + 3 energies
Phase changes
Solid to gas is sublimation, gas to solid is deposition
When vapor pressure equals the surrounding atmospheric pressure, the liquid boils
Enthalpy of fusion - energy to melt a solid; heat of fusion - heat given off by a substance when it freezes
Enthalpy of vaporization - energy to turn a liquid into a gas; heat of vaporization - heat given off by a substance condensing
IMF is stronger for a liquid than a gas, and the stronger IMF is more stable, therefore going from a gas to a liquid or a liquid to a solid releases energy (exothermic)
As heat is added to a substance, the temperature of the substance can increase OR it can change phases, but not both at once
When a substance is changing phases, the temperature of the substance remains constant
Calorimetry
Specific heat - amount of heat required to raise the temperature of one gram of a substance by one degree C/K
Large specific heat - can absorb much heat without a significant temperature change
Low specific heat - quickly changes temperature
q = mcΔT
q1 = q2 for mixtures
Calorimetry - measurement of heat changes during chemical reactions
Find J from q, find mol from stoich, divide the two to find delta H
Delta H measured in J/mol
Heating curves
For problems where a solid completely melts or the like, add q from MCAT to (moles) * (heat of fusion) for the total heat required for process to occur
VII. Equilibrium
Keq
Reaction is at equilibrium when all concentrations stop changing
Reaction does not stop - rate of forward and reverse reactions become equal
All concentrations do NOT sum to initial concentration of reactants
Keq expression: For the reaction aA + bB -> cC + dD: Keq = ([C]^c [D]^d) / ([A]^a [B]^b)
[A], etc. are molar concentrations/partial pressures at equilibrium
Products in numerator, reactants in denominator
Coefficients in balanced equation become exponents in equilibrium expression
Only gaseous and aqueous species are included in the expression
Keq has no units
K>1 favors forward rxn; K<1 favors reverse rxn
Different equilibrium constants
Kc for molar concentrations
Kp for partial pressures
Ksp is solubility product (no denominator because reactants are solids)
Ka is acid dissociation constant for weak acids
Kb is base dissociation constant for weak bases
Kw describes the ionization of water (Kw = 1*10^-14)
Manipulating Keq
Keq for a flipped reaction is the reciprocal of Keq for initial rxn
Keq for a reaction multiplied by a coefficient is the initial Keq to the power of the coefficient
Keq for two reactions added together is their respective initial Keq values multiplied together
Le Chatelier’s principle
Increasing concentration of reactants shifts rxn to favor products (forward) and vice versa
Increasing pressure increases partial pressure of all gases in container and shifts rxn to side with fewer gas molecules
Increasing volume decreases pressure and vice versa
Adding a non-reacting gas to a non-rigid container causes the volume to increase while not changing total pressure
Adding a non-reacting gas to a rigid container would increase the total pressure of the container and not affect the partial pressures of other species - no reaction shift occurs
Increasing temperature in an endothermic reaction shifts the rxn to favor products (forward); increasing temperature in an exothermic reaction shifts the rxn to favor reactants (reverse)
Treat “heat” as a reactant (endothermic) or product (exothermic) to see shifts like with concentration change
Diluting aqueous equilibriums shifts the rxn to favor the side with more aqueous species; removing water (evaporation) shifts the rxn to favor the side with less aqueous species
Shifts caused by concentration/pressure are temporary shifts and do not change Keq; shifts caused by temperature permanently affects Keq and ratio of products to reactants since it adds/removes energy from the system
Reaction quotient Q
Q can be calculated at any point with current concentrations/pressures; Keq can only be calculated with equilibrium values
For the reaction aA + bB -> cC + dD: Q = ([C]^c [D]^d) / ([A]^a [B]^b)
[A], etc. are initial molar concentrations or partial pressures
If QK, rxn shifts left; if Q=K, rxn is at equilibrium
Solubility
A salt is considered soluble if more than 1g can be dissolved in 100mL of water
Soluble salts are assumed to dissociate completely in aqueous solutions
Most solids become more soluble in a liquid as temp increases
Solubility product Ksp
For the reaction AaBb(s) ⇄ aA^b+(aq) + bB^a-(aq): Ksp = [A^b+]^a * [B^a-]^b
Molar solubility is determined by subbing x, 2x, 3x, etc. in for concentrations in Ksp expression (x if coefficient is 1 in balanced reaction, 2x if coefficient is 2, etc.)
Molar solubility of a salt is equal to the concentration of any ion that occurs in a 1:1 ratio with the salt
Molar solubility typically increases with temperature since there is more energy available to force water molecules apart to make room for solute ions
Common ion effect
Newly added ions from a separate solution affect equilibrium of initial solution if some elements are present in both, even though newly added ions did not come from the initial compound
ex. Adding NaCl to AgCl affects Cl which affects AgCl equilibrium
VIII. Acids and Bases
pH
Formulas
pH = -log([H+])
pOH = -log([OH-])
pKa = -log(Ka)
pKb = -log(Kb)
pKw = -log(Kw)
[H+] = [OH-] => neutral, pH = 7
[H+] > [OH-] => acidic, pH < 7
[H+] < [OH-] => basic, pH > 7
Increasing pH means decreasing [H+] (less acidic solution) and vice versa
Strong acids
Strong acids dissociate completely in water (rxn goes to completion); no equilibrium, eq constant, or dissociation constant
Important strong acids/bases
No tendency for reverse rxn to occur -> conjugate base of a strong acid is very weak
pH of strong acid solution can be found directly from [H+] since it dissociates completely
Best conductors of electricity
Weak acids
Weak acid + water causes a small fraction of its molecules to dissociate into H+ and A- (conjugate base) ions
Ka and Kb are measures of the strengths of strong/weak acids - equilibrium constants specific to acids/bases
Acid dissociation constant Ka = [H+]*[A-]/[HA]
Base dissociation constant Kb = [HB+]*[OH-]/[B]
Greater Ka means a greater extent of dissociation and a stronger acid
Greater Kb means a stronger base; base is not dissociating but rather accepting a proton from an acid
Set up RICE table w/ values of x for gained/lost concentration to solve for [H+] and pH from Ka or vice versa
Acid Strength
Percent dissociation
The more H+ ions an acid can donate, the stronger the acid is
Lower concentration -> higher percent dissociation; higher concentration will lead to more of the conjugate base, making it easier for the reverse rxn to take place -> more HA present in solution and less H3O+ ions (lower percent dissociation)
Percent ionization: [H3O+]/[HA] * 100
Acid/base structure
H is written in front of acids even if H is contained in the conjugate base because that H is attached to an (usually O) atom at the end of the molecule
H in a hydroxyl group (-OH) are dissociable due to O being more electronegative than H
H bonded to C is almost never dissociable
Solubility
Hydroxides dissolve well in solutions with low pH (more H+ ions to react
with OH- and speed rxn along)
Polyprotic acids
Acids that can give up more than one hydrogen ion (ex. H3PO4)
More willing to give up first proton than others (after 1st, resulting negative charge attracts remaining protons more strongly)
H3PO4 is a stronger acid than H2PO4-, HPO42-, etc.
Amount of each succeeding acid decreases: [H3PO4] > [H2PO4-] > [HPO42-] > [PO43-]
Kw
The equilibrium constant of water due to the following reaction: Kw = [H3O+][OH-] = [H+][OH-] = 1.0*10^-14 at 25 C for any aqueous solution
pH + pOH = 14
Kw = 110^-14 = KaKb
pKa + pKb = 14
Knowing Ka for a weak acid, Kb can be found for its conjugate base
pH is not limited to a 0-14 scale - very rarely is pH >14 or <0, but it does occur at high concentrations
Increase in temperature increases Kw (dissociation of water is endothermic) so pKw and pH decrease
Neutralization reactions
When an acid and base mix, the acid donates protons to the base in a neutralization rxn
Strong acid + strong base
Both substances dissociate completely
Net ionic is always the creation of water: H+(aq) + OH-(aq) ⇄ H2O(l)
All other ions are spectator ions
Strong acid + weak base
Strong acid (which dissociates completely) will donate a proton to the weak base
Product is conjugate acid of weak base
Ex. HCl + NH3: Net ionic is H+(aq) + NH3(aq) ⇄ NH4+(aq)
Weak acid + strong base
Strong base will accept protons from weak acid
Products are conjugate base of weak acid and water
Ex. HC2H3O2 + NaOH: Net ionic is HC2H3O2(aq) + OH-(aq) ⇄ C2H3O2-(aq) + H2O(l)
Weak acid + weak base
Simple proton transfer reaction - acid gives protons to base
Ex. HC2H3O2 + NH3: Net ionic is HC2H3O2(aq) + NH3(aq) ⇄ C2H3O2-(aq) + NH4+(aq)
Buffers
Solution with a very stable pH; acid/base can be added to a buffer solution without greatly affecting pH; gain/loss of water also does not change pH
Buffers are created by placing large amounts of a weak acid/base into a solution with its conjugate (salt)
Weak acid and conjugate base remain in solution together without neutralizing each other
Presence of the conjugate pair makes the buffer effective
If enough strong acid/base is added that all of the acid or conjugate base is reacted, the buffer breaks
Higher concentrations of the conjugate pair resist pH change (better buffers) better than lower concentrations
Henderson-Hasselbalch
When concentrations of acid and conjugate base in a solution are the same, pH=pKa and pOH=pKb
Choosing an acid for a buffer solution requires choosing an acid with a pKa close to the desired pH

(almost equal amounts of acid and conjugate base; makes buffer flexible in neutralizing both added H+ and OH-)

Buffers cannot be created from a very strong acid and its conjugate base because the conjugate base will be very weak and will not readily accept protons
Indicators
Weak acids which change colors in certain pH ranges due to LeChatelier’s principle
HIn ⇄ H+ + In-
Ka = [H+][In-]/[HIn]
Protonated HIn state must be a different color from deprotonated In- state
Acidic environment causes excess H+ to drive equilibrium to the left (color of HIn); basic environment causes excess OH- to react with H+ from indicator and drive reaction right (color of In-)
Color change occurs when [HIn] = [In-]; or pH = pKa
Choose an indicator whose pKa matches the pH at the titration’s equivalence point
Titration
Neutralization reactions are performed by titration, where a base of known concentration is slowly added to an acid or vice versa
Titration curves
Midpoint also called half equivalence point occurs when [HA] = [A-] (pH = pKa)
Equivalence point occurs when just enough base has been added to neutralize all the acid initially present (equimolar)
HA, A- present before midpoint; A- at midpoint, OH- after midpoint
X. Laboratory Overview
Weighing hot objects on a scale creates convection currents, making object appear lighter than it truly is
Not rinsing a buret in a titration leads to it being diluted

Figure: Molecular geometry (VSEPR) ####

[Image: Molecular geometry (VSEPR) diagram showing various shapes such as linear, trigonal planar, tetrahedral, etc.]

III. Intermolecular Forces and Properties (continued)
Polynomial (assorted) details shown in figure
Relationship of geometry to polarity and IMF

Common Polyatomic Ions Table ####

Common Polyatomic Ions
C2H3O2- Acetate
NH4+ Ammonium
CO3^2- Carbonate
ClO3- Chlorate
ClO4- Perchlorate
CrO4^2- Chromate
CN- Cyanide
Cr2O7^2- Dichromate
HCO3- Bicarbonate
HSO4- Bisulfate
HSO3- Bisulfite
OH- Hydroxide
ClO- Hypochlorite
NO3- Nitrate
NO2- Nitrite
C2O4^2- Oxalate
MnO4- Permanganate
PO4^3- Phosphate
SO4^2- Sulfate
SO3^2- Sulfite

IV. Chemical Reactions
Types of reactions
Synthesis: everything combines to form one compound
Decomposition: one compound + heat is split into multiple elements/compounds
Acid-base rxn: Acid + base -> water + salt
Oxidation-reduction (redox) rxn: changes the oxidation state of some species
Combustion: substance w/ H and C + O2 -> CO2 + H2O
Precipitation: aqueous solutions -> insoluble salt (+ more aq sometimes)
Can be written as net ionic - Those free ions not in net ionic are spectator ions
Solubility rules
Alkali metal cations or ammonium (NH4+) cations are ALWAYS soluble
Compounds with a nitrate (NO3-) anion are ALWAYS soluble
Common polyatomic ions
Calculations
Percent error: 100 * abs(experimental - expected)/(expected)
Combustion analysis - use law of conservation of mass
Gravimetric analysis - determine identity by comparing mass percentages
Oxidation states
Neutral atoms 0
Monoatomic ions equal to ion charge
Oxygen -2 (except in H2O2 where it is -1)
Hydrogen +1 with nonmetals; -1 with metals
In absence of oxygen, most electronegative element takes usual charge
C, N, S, P vary often
Redox reactions
Write 2 half-reactions; balance O with H2O; balance H with H+; combine
ACIDIC: stop here
BASIC: add OH- to both sides to neutralize H+; cancel H2O
Acids and bases (briefly)
Color change signals end of titration
Acids donate H+, bases donate electrons
Water is amphoteric
V. Kinetics
Rate law: Rate = k [A]^x [B]^y [C]^z
Determine x, y, z via experiments
Keq = forward/backward rate constants
Units: rate M/s; k units depend on order
Reaction stoichiometry relation for rates
Orders: zero, first, second order rules; half-life
Collision theory
Activation energy Ea; orientation requirements; temperature and concentration effects
Reaction energy profile
Reaction mechanisms
Catalysts
VI. Thermodynamics
Temperature/heat
Temperature = average kinetic energy
Heat = energy flow between systems
First law
Bond formation/releases energy
Exothermic vs Endothermic
Energy diagrams
Enthalpy
Formation; Hess’s law; bond energy; enthalpy of solution
Hydration energy; Coulombic energy
Phase changes
Sublimation/Deposition; boiling point; fusion/vaporization; IMF role
Calorimetry
Specific heat; q = mcΔT; q1 = q2; calorimetry for reactions; delta H
Heating curves
VII. Equilibrium
Keq; rate concepts; Le Chatelier; Q vs Keq
K values: Kc, Kp, Ksp, Ka, Kb, Kw
Manipulating Keq
Le Chatelier's principle specifics
Solubility and common ion effects
VIII. Acids and Bases
pH and related formulas
Strong vs weak acids; acid strength; Ka, Kb; RICE table
Kw relations; pH/pKw; buffer concepts; Henderson-Hasselbalch
Buffers and indicators; titration concepts; equivalence and midpoint
IX. Applications of Thermodynamics
Entropy; standard entropies; delta S; Gibbs free energy; Delta G; spontaneity
Galvanic/voltaic cells; anode/cathode; electron flow; salt bridge
Nernst equation; non-standard conditions; cell potential
Electrolytic cells; calculation of metal deposition
Voltage and favorability
X. Laboratory Overview
Practical notes: weighing hot objects; rinsing burets

The content above mirrors typical AP Biology topics with detailed subpoints and examples as presented in the source document, including electron configuration, bonding types, thermodynamics, equilibria, acids/bases, and laboratory principles. All sections, subsections, and data points are preserved to enable reconstruction of the original document from this JSON export.

Figure: Molecular geometry (VSEPR) ####

Image reference: Molecular geometry (VSEPR) diagram showing various shapes like linear, trigonal planar, tetrahedral, trigonal pyramidal, square planar, etc. {#figure-molecular-geometry}

Image details

Alt text: “Molecular geometry (VSEPR) diagram illustrating common geometries and steric numbers”
Caption: “Figure: Molecular geometry examples according to VSEPR theory.”

Images

[Image] url: images/molecular-geometry-vsepr-diagram.jpg; alt: Molecular geometry diagram per VSEPR; caption: Figure: Molecular geometry (VSEPR) illustration

AP Biology Review

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Periodic table

Alkali metals, alkaline earth metals, transition metals, halogens, noble gases

Mass number = P + N

Isotopes - atoms of an element with different numbers of neutrons

Average atomic mass from weighted average of isotope mass and relative abundance (frequency)

Moles

PV = nRT

Avogadro’s number 6.022*10^23

AT STP (1 atm, 273K), 22.4 L/mol

Molarity M = moles/L

Percent composition - divide the mass of each element/compound by the total molar mass of the substance

Empirical formula is simplest ratio, molecular formula is actual formula for substance

Energy

Electron potential energy increases with distance from nucleus

Electron energy is quantized - can only exist at specific energy levels at specific intervals, not in between

Coulomb’s law: F = kq1q2/(r^2) where F is electrostatic force

Atoms absorb energy in the form of electromagnetic radiation as electrons jump to higher energy levels; when electrons drop levels (closer), atoms give off energy

Photoelectron spectroscopy

energy measured in electronvolts (eV)

Incoming radiation energy = binding energy + kinetic energy of the ejected electron

Electrons that are further away from nucleus require less energy to eject, thus will move faster

Photoelectron spectrum

Each section of peaks represents a different energy level (1, 2, 3, etc.)

Subshells within each energy level (shape of space electron can be found in orbiting nucleus) are represented by the peaks (1s, 2s, 2p, etc.)

s(2) - first subshell, p(6) - second subshell

Height of peaks determines number of electrons in subshell (ex. Peak of p subshell in energy level 2 will be 3x that of s subshell)

Electron configuration

Electron configuration - spdf - shorthand with noble gas first

Configuration rules

Aufbau principle - electrons fill lowest energy subshells available first

Pauli exclusion principle - 2 electrons in same orbital cannot have same spin

Hund’s rule - Electrons occupy empty subshells first

Zn +2, Ag +1, Al +3, Cd +2, most other transition metal charges vary

Periodic trends

Electrons are more attracted if they are closer to the nucleus, or if there are more protons

Electrons are repelled by other electrons - if there are electrons b/w the valence electrons and nucleus, the e- will be less attracted (shielding)

Completed shells are very stable, completed subshells are also stable; atoms will add/subtract valence electrons to complete their shell

INCREASING: atomic radius down left; ionization energy up right; electronegativity up right

Ionization energy - energy required to remove an electron from an atom

Electronegativity - how strongly the nucleus of an atom attracts electrons of other atoms in a bond

Electron affinity - energy change that occurs when an electron is added to an atom in the gas state (usually exothermic - energy is released)

II. Molecular and Ionic Compound Structure and Properties

Bonds

Atoms are more stable with full valence shells

Ionic bonds

Cation gives up electrons completely

Electrostatic attractions in a lattice structure

Metals and nonmetals (salts)

Coulomb’s law - greater charge leads to a greater bond/lattice energy (higher melting point)

If both have equal charges, smaller radius will have greater coulombic attraction

Ionic solid - electrons do not move around lattice; ionic solids are poor conductors of electricity; ionic liquids conduct electricity because ions are free to move around, though e- are still localized around particular atoms

Metallic bonds

Sea of electrons model - positively charged core is stationary while valence electrons are very mobile

Metals bond to form alloys - interstitial alloy w/ metals of different radii; substitutional alloy w/ metals of similar radii

Molecular covalent bonds

2 atoms share electrons - both atoms achieve complete outer shells

2 nonmetals

Creates molecules - 2+ atoms covalently bonded together

Single has 1 sigma bond - order 1, longest length, least energy; double has 1 sigma and 1 pi bond - order 2, int. length, int. energy; triple has 1 sigma and 2 pi bonds - order 3, shortest length, greatest bond energy

Bond forms when potential energy is at minimal level

Too close - potential energy is too high due to repulsive forces

Too far - potential energy is near 0 because attractive forces are very weak

Minimul PE occurs when repulsive and attractive forces are balanced

Network covalent bonds - lattice of covalent bonds - poor conductors, high melting and boiling points

Conductivity

Conductivity of different substances in different phases

Solid Aqueous Liquid Gas

Ionic No Yes Yes No

Molecular Covalent No No No No

Network Covalent No N/A No No

Metallic Yes N/A Yes No

Lewis dot structures

Resonance - for bond order calculations, average together all possible orders of a specific bond

BORON (B) is stable with 6 electrons - only one that does not need a full octet

Expanded octets - any atom of an element from n=3 or greater (those with a d subshell) can have [8,12] valence electrons on center atom

Noble gases form bonds by filling empty d orbital with electrons

Formal charge - number of valence electrons minus assigned electrons (1 e- for each line “shared” bond) - 0 for neutral molecules

Molecular geometry (VSEPR)

Double and triple bonds have more repulsive strength than single bonds - occupy more space

Lone electron pairs have more repulsive strength than bonding pairs, so molecules with lone pairs will have slightly reduced angles between terminal atoms

Hybridization - how many atoms are attached (sp, sp2, sp3, sp3d, etc.)

III. Intermolecular Forces and Properties

Polarity

Covalent bond where electrons are unequally shared - polar covalent

Dipoles are caused by polar covalent bonds - pair of opposite electric charges separated by some distance, like partial charges on atoms in a polar covalent bond

If 2 identical atoms bond (ex. Cl-Cl) the electrons are equally shared, creating a nonpolar covalent bond with no dipole

Bonds can be polar; so can molecules depending on the molecular geometry (and polarity of bonds - secondary)

In polar molecules, more electronegative atoms will gain negative partial charge; Usually central atom will be positive - exception is hydrogen (terminal), which is usually positive since it has less electronegativity

Intermolecular forces

Forces b/w molecules in a covalently bonded substance - need to be broken apart for covalent substances to change phases

Changing phase: ionic substances break bonds b/w individual ions; covalent substances keep bonds inside a molecule in place but break bonds b/w molecules

Dipole-dipole forces

Polar molecules - positive end of one molecule is attracted to negative end of another molecule

Greater polarity -> greater dipole dipole attraction -> larger dipole moment -> higher melting/boiling points

Relatively weak overall - melt and boil at low temps

Hydrogen bonds

Special type of dipole-dipole attraction where positively charged hydrogen end of a molecule is attracted to negatively charged end of another molecule containing an extremely electronegative element (F, O, N)

Much stronger than normal dipole-dipole forces since a hydrogen atom sharing/giving up its lone e- to a bond is left w/ no shielding

Higher melting/boiling points than substances held together only by other types of IMF

London dispersion forces

All molecules - very weak attractions due to random motion of electrons on atoms within molecules (instantaneous polarity)

Molecules w/ more e- experience greater LDF (more random motion)

Higher molar mass usually means greater LDF (as mass increases, e- increases for the molecule to remain electrically neutral)

IMF strength

Ionic substances are generally solids at room temp - melting them requires lattice bonds to be broken - energy determined by Coulombic attraction

Covalent substances (liquids at room temp) boil when IMF are broken; for molecules of similar size, from strongest to weakest: hydrogen bonds, permanent dipoles, LDF (temporary dipoles - greater for larger molecules)

Melting/boiling points of covalent substances are LOWER than for ionic substances

Bonding/Phases

Substances w/ weak IMF (LDF) tend to be gases at room temp (N2); substances w/ strong IMF (hydrogen bonds) tend to be liquids at room temp (H2O)

Ionic substances do not experience IMF - since ionic bonds are stronger than IMF, ionic substances are usually solids at room temp

Vapor pressure

Molecules in a liquid are in constant motion - if they hit the surface of the liquid with enough kinetic energy, they can escape the IMF holding them to other molecules and transition into the gas phase

Vaporization (NOT boiling) - no outside energy is added

Temperature and vapor pressure are directly proportional

At the same temp, vapor pressure is dependent on strength of IMF (stronger IMF, lower vapor pressure)

Solution separation

Solutes and solvents - like dissolves like

Paper chromatography

Piece of filter paper with substance on the bottom is dipped in water

More polar components of substance travel further up the filter paper with the polar water

Distance substance travels up the paper measured by retention/retardation factor Rf = (distance traveled by solute - substance being separated)/(distance traveled by solvent front - water)

Stronger attraction - larger Rf

Column chromatography

Column is packed with a stationary substance

separable solution (analyte) is injected, adhering to stationary phase

another solution (eluent - liquid/gas) is injected into column

more attracted analyte molecules will move through faster and leave column first

Distillation

Takes advantage of different boiling points of substances by boiling a mixture at an intermediate point

Kinetic molecular theory

Kinetic energy of a single gas molecule: KE = ½ mv^2

Average kinetic energy of a gas depends on the temperature (directly proportional), not the identity of the gas

Ideal gases have insignificant volume of molecules, no forces of attraction b/w molecules, and are in constant motion without losing KE

Deviations occur at low temperatures or high pressures

Volume becomes significant

Real gases have attractions and lower pressure than predicted

Maxwell-boltzman diagrams

Higher temp -> greater KE -> greater range of velocity

Smaller masses, greater velocities to have same KE

Effusion

Rate at which a gas escapes from a container through microscopic holes

High to low pressure

Greater speed, greater temp, greater rate of effusion

If at same temp, gas w/ lower molar mass will effuse first

Equations

Ideal gas equation: PV = nRT

R = 0.0821

Combined gas law: P1V1/T1 = P2V2/T2

Dalton’s law: P(total) = Pa + Pb + Pc + …

Partial pressure: Pa = P(total)*(moles of gas A)/(total moles of gas)

Density: D = m/V

From ideal gas law: Molar mass = DRT/P

Electromagnetic spectrum

E = hv; E = energy change; h = Planck’s constant 6.626*10^-34; v = frequency

C = lambda v; C = speed of light 2.99810^8; v = frequency; lambda = wavelength

Beer’s law: A = abc

A = absorbance; a = molar absorptivity; b = path length; c = concentration

Colorimetry - direct relationship between concentration and absorbance

IV. Chemical Reactions

Types of reactions

Synthesis: everything combines to form one compound

Decomposition: one compound + heat is split into multiple elements/compounds

Acid-base rxn: Acid + base -> water + salt

Oxidation-reduction (redox) rxn: changes the oxidation state of some species

Combustion: substance w/ H and C + O2 -> CO2 + H2O

Precipitation: aqueous solutions -> insoluble salt (+ more aq sometimes)

Can be written as net ionic - Those free ions not in net ionic are spectator ions

Solubility rules

Alkali metal cations or ammonium (NH4+) cations are ALWAYS soluble

Compounds with a nitrate (NO3-) anion are ALWAYS soluble

Common polyatomic ions

Calculations

Percent error: 100 * abs(experimental - expected)/(expected)

Combustion analysis - use law of conservation of mass (if x g of CO2 is produced, find g of C which will be starting amt)

Gravimetric analysis - when asked to determine the identity of a certain compound, find g of component produced, then use mass percent (g found / total sample mass) and compare to mass percent of options (molar mass of component / molar mass of entire compound)

Oxidation states

Neutral atoms not bonded to other atoms have an oxidation state of 0

Monoatomic ions have an oxidation state equal to the charge on that ion (ex. Zn2+ will be +2)

Oxygen is -2 (EXCEPTION: in hydrogen peroxide, H2O2, O is -1)

Hydrogen is +1 w/ nonmetals, -1 with metals

In absence of oxygen, most electronegative element in a compound will take an oxidation state equal to its usual charge (ex. F is -1 in CF4)

IF none of the above rules apply, determine the oxidation state by adding up all elements’ oxidation states to 0/charge on ion

C, N, S, P frequently vary oxidation states (low electronegativity)

Redox reactions

Write full rxn as 2 half reactions (oxidation and reduction; OIL RIG)

Add H2O to compensate for oxygen on one side

Add H+ to compensate for H from H2O on other side

Balance 2 half rxns to have the same number of electrons and add them together to produce one complete reaction

ACIDIC: stop here

BASIC: Add OH- to both sides - enough for all H+ on one side to be converted to H2O; then cancel out H2O so it only remains on one side

Acids and bases (briefly)

Color change signals the end of a titration (can be redox or acid/base)

Acids are capable of donating protons (H+); bases are capable of donating electrons

Water can act as an acid or base - amphoteric

V. Kinetics

Rate law

Rate = k [A]^x [B]^y [C]^z

Can calculate x, y, z via a table from (concentration factor)^x = (rate factor)

K is only dependent on temperature (always increases w/ T)

Keq = K1 (rate constant of forward rxn) / K2 (rate constant of reverse rxn)

K calculated by dividing any rate in table by the concentrations to their respective powers

Units for rate are M/s, units for conc are M -> calculate units for k from there

If A + 2B + C -> D; rate of formation of D = rate of disappearance of A and C = 0.5*rate of disappearance of B

Orders

Zero-order: Rate = k ; Concentration vs time has slope -k

First-order: Rate = k[A]; ln[A] vs time has slope -k; ln[A]t = -kt + ln[A]0

Second-order: Rate = k[A]^2; 1/[A] vs time has slope k; 1/[A]t = kt + 1/[A]0

Half-life: First order reactions only have a constant half life; t1/2 = ln(2)/k = 0.693/k

Collision theory

Chemical reactions occur because reactants are constantly moving and colliding with one another

When reactants collide with sufficient energy (activation energy Ea), a reaction occurs

Gaseous/aqueous: increased concentration increases rate of reaction

Stirring increases reaction rate for heterogeneous mixtures

Greater temp increases reaction rate

Reactions only occur if reactants collide with correct orientation

Reaction energy profile

Reaction mechanisms

Intermediates exist as species produced in a mechanism and consumed later

Adding up steps and canceling gives the overall equation

Elementary steps w/ 2 reactants are bimolecular; w/ 1 reactant are unimolecular

Speed is determined by slow step (rate-determining step)

Consistency is determined by slow step and those leading up to it

Slow step has highest activation energy

Catalysts

Catalysts increase rate without being consumed; do not appear in balanced equation

In a mechanism, catalysts enter first, then exit

Types: surface catalysis, enzyme catalysis, acid-base catalysis

VI. Thermodynamics

Temperature/heat

Temperature is the average amount of kinetic energy

Heat is the energy flow between two substances at different temperatures

First law of thermodynamics: energy can be neither created nor destroyed

When bonds form, energy is released; when bonds are broken, energy is absorbed

Exothermic - energy transferred from system to surroundings (delta H negative)

Endothermic - energy transferred from surroundings into system (delta H positive)

Energy diagrams

Enthalpy

Enthalpy of formation

Change in energy when one mole of a compound is formed from its component pure elements under standard conditions (25C/298K)

Delta Hf = delta Hf for products - delta Hf for reactants

If delta Hf is negative, energy is released when the compound is formed, so the product is more stable (exothermic)

If delta Hf is positive, energy is absorbed when the compound is formed, so the product is less stable than its constituent elements (endothermic)

Heat of formation is 0 when the pure element is in its standard state (ex. H2(g) or F2(g))

Bond energy

Delta H (J) = bond energies of reactants - bond energies of products

Multiply bond energies for each bond by the coefficient

Hess’s law

Finding delta H for the overall reaction from knowing delta H for the steps of the reaction

Flipping the equation flips the sign of delta H

Multiplying/dividing the equation by a coefficient multiplies/divides delta H by that coefficient

Adding/subtracting equations adds/subtracts their delta H values

Enthalpy of solution

Ionic substances dissolving in water

1: Breaking of solute bonds - energy required is equal to the lattice energy (positive delta H)

2: Separation of solvent molecules - water molecules must spread out to make room for the solute ions (positive delta H)

3: Creation new attractions - free floating ions are attracted to the dipoles of water molecules (negative delta H)

Hydration energy = step 2 + step 3 energies

Coulombic energy - increases with charge magnitude, decreases as size increases

Enthalpy of solution = step 1 + 2 + 3 energies

Phase changes

Solid to gas is sublimation, gas to solid is deposition

When vapor pressure equals the surrounding atmospheric pressure, the liquid boils

Enthalpy of fusion - energy to melt a solid; heat of fusion - heat given off by a substance when it freezes

Enthalpy of vaporization - energy to turn a liquid into a gas; heat of vaporization - heat given off by a substance condensing

IMF is stronger for a liquid than a gas, and the stronger IMF is more stable, therefore going from a gas to a liquid or a liquid to a solid releases energy (exothermic)

As heat is added to a substance, the temperature of the substance can increase OR it can change phases, but not both at once

When a substance is changing phases, the temperature of the substance remains constant

Calorimetry

Specific heat - amount of heat required to raise the temperature of one gram of a substance by one degree C/K

Large specific heat - can absorb much heat without a significant temperature change

Low specific heat - quickly changes temperature

q = mcΔT

q1 = q2 for mixtures

Calorimetry - measurement of heat changes during chemical reactions

Find J from q, find mol from stoich, divide the two to find delta H

Delta H measured in J/mol

Heating curves

For problems where a solid completely melts or the like, add q from MCAT to (moles) * (heat of fusion) for the total heat required for process to occur

VII. Equilibrium

Keq

Reaction is at equilibrium when all concentrations stop changing

Reaction does not stop - rate of forward and reverse reactions become equal

All concentrations do NOT sum to initial concentration of reactants

Keq expression: For the reaction aA + bB -> cC + dD: Keq = ([C]^c [D]^d) / ([A]^a [B]^b)

[A], etc. are molar concentrations/partial pressures at equilibrium

Products in numerator, reactants in denominator

Coefficients in balanced equation become exponents in equilibrium expression

Only gaseous and aqueous species are included in the expression

Keq has no units

K>1 favors forward rxn; K<1 favors reverse rxn

Different equilibrium constants

Kc for molar concentrations

Kp for partial pressures

Ksp is solubility product (no denominator because reactants are solids)

Ka is acid dissociation constant for weak acids

Kb is base dissociation constant for weak bases

Kw describes the ionization of water (Kw = 1*10^-14)

Manipulating Keq

Keq for a flipped reaction is the reciprocal of Keq for initial rxn

Keq for a reaction multiplied by a coefficient is the initial Keq to the power of the coefficient

Keq for two reactions added together is their respective initial Keq values multiplied together

Le Chatelier’s principle

Increasing concentration of reactants shifts rxn to favor products (forward) and vice versa

Increasing pressure increases partial pressure of all gases in container and shifts rxn to side with fewer gas molecules

Increasing volume decreases pressure and vice versa

Adding a non-reacting gas to a non-rigid container causes the volume to increase while not changing total pressure

Adding a non-reacting gas to a rigid container would increase the total pressure of the container and not affect the partial pressures of other species - no reaction shift occurs

Increasing temperature in an endothermic reaction shifts the rxn to favor products (forward); increasing temperature in an exothermic reaction shifts the rxn to favor reactants (reverse)

Treat “heat” as a reactant (endothermic) or product (exothermic) to see shifts like with concentration change

Diluting aqueous equilibriums shifts the rxn to favor the side with more aqueous species; removing water (evaporation) shifts the rxn to favor the side with less aqueous species

Shifts caused by concentration/pressure are temporary shifts and do not change Keq; shifts caused by temperature permanently affects Keq and ratio of products to reactants since it adds/removes energy from the system

Reaction quotient Q

Q can be calculated at any point with current concentrations/pressures; Keq can only be calculated with equilibrium values

For the reaction aA + bB -> cC + dD: Q = ([C]^c [D]^d) / ([A]^a [B]^b)

[A], etc. are initial molar concentrations or partial pressures

If QK, rxn shifts left; if Q=K, rxn is at equilibrium

Solubility

A salt is considered soluble if more than 1g can be dissolved in 100mL of water

Soluble salts are assumed to dissociate completely in aqueous solutions

Most solids become more soluble in a liquid as temp increases

Solubility product Ksp

For the reaction AaBb(s) ⇄ aA^b+(aq) + bB^a-(aq): Ksp = [A^b+]^a * [B^a-]^b

Molar solubility is determined by subbing x, 2x, 3x, etc. in for concentrations in Ksp expression (x if coefficient is 1 in balanced reaction, 2x if coefficient is 2, etc.)

Molar solubility of a salt is equal to the concentration of any ion that occurs in a 1:1 ratio with the salt

Molar solubility typically increases with temperature since there is more energy available to force water molecules apart to make room for solute ions

Common ion effect

Newly added ions from a separate solution affect equilibrium of initial solution if some elements are present in both, even though newly added ions did not come from the initial compound

ex. Adding NaCl to AgCl affects Cl which affects AgCl equilibrium

VIII. Acids and Bases

Formulas

pH = -log([H+])

pOH = -log([OH-])

pKa = -log(Ka)

pKb = -log(Kb)

pKw = -log(Kw)

[H+] = [OH-] => neutral, pH = 7

[H+] > [OH-] => acidic, pH < 7

[H+] < [OH-] => basic, pH > 7

Increasing pH means decreasing [H+] (less acidic solution) and vice versa

Strong acids

Strong acids dissociate completely in water (rxn goes to completion); no equilibrium, eq constant, or dissociation constant

Important strong acids/bases

No tendency for reverse rxn to occur -> conjugate base of a strong acid is very weak

pH of strong acid solution can be found directly from [H+] since it dissociates completely

Best conductors of electricity

Weak acids

Weak acid + water causes a small fraction of its molecules to dissociate into H+ and A- (conjugate base) ions

Ka and Kb are measures of the strengths of strong/weak acids - equilibrium constants specific to acids/bases

Acid dissociation constant Ka = [H+]*[A-]/[HA]

Base dissociation constant Kb = [HB+]*[OH-]/[B]

Greater Ka means a greater extent of dissociation and a stronger acid

Greater Kb means a stronger base; base is not dissociating but rather accepting a proton from an acid

Set up RICE table w/ values of x for gained/lost concentration to solve for [H+] and pH from Ka or vice versa

Acid Strength

Percent dissociation

The more H+ ions an acid can donate, the stronger the acid is

Lower concentration -> higher percent dissociation; higher concentration will lead to more of the conjugate base, making it easier for the reverse rxn to take place -> more HA present in solution and less H3O+ ions (lower percent dissociation)

Percent ionization: [H3O+]/[HA] * 100

Acid/base structure

H is written in front of acids even if H is contained in the conjugate base because that H is attached to an (usually O) atom at the end of the molecule

H in a hydroxyl group (-OH) are dissociable due to O being more electronegative than H

H bonded to C is almost never dissociable

Solubility

Hydroxides dissolve well in solutions with low pH (more H+ ions to react

with OH- and speed rxn along)

Polyprotic acids

Acids that can give up more than one hydrogen ion (ex. H3PO4)

More willing to give up first proton than others (after 1st, resulting negative charge attracts remaining protons more strongly)

H3PO4 is a stronger acid than H2PO4-, HPO42-, etc.

Amount of each succeeding acid decreases: [H3PO4] > [H2PO4-] > [HPO42-] > [PO43-]

The equilibrium constant of water due to the following reaction: Kw = [H3O+][OH-] = [H+][OH-] = 1.0*10^-14 at 25 C for any aqueous solution

pH + pOH = 14

Kw = 110^-14 = KaKb

pKa + pKb = 14

Knowing Ka for a weak acid, Kb can be found for its conjugate base

pH is not limited to a 0-14 scale - very rarely is pH >14 or <0, but it does occur at high concentrations

Increase in temperature increases Kw (dissociation of water is endothermic) so pKw and pH decrease

Neutralization reactions

When an acid and base mix, the acid donates protons to the base in a neutralization rxn

Strong acid + strong base

Both substances dissociate completely

Net ionic is always the creation of water: H+(aq) + OH-(aq) ⇄ H2O(l)

All other ions are spectator ions

Strong acid + weak base

Strong acid (which dissociates completely) will donate a proton to the weak base

Product is conjugate acid of weak base

Ex. HCl + NH3: Net ionic is H+(aq) + NH3(aq) ⇄ NH4+(aq)

Weak acid + strong base

Strong base will accept protons from weak acid

Products are conjugate base of weak acid and water

Ex. HC2H3O2 + NaOH: Net ionic is HC2H3O2(aq) + OH-(aq) ⇄ C2H3O2-(aq) + H2O(l)

Weak acid + weak base

Simple proton transfer reaction - acid gives protons to base

Ex. HC2H3O2 + NH3: Net ionic is HC2H3O2(aq) + NH3(aq) ⇄ C2H3O2-(aq) + NH4+(aq)

Buffers

Solution with a very stable pH; acid/base can be added to a buffer solution without greatly affecting pH; gain/loss of water also does not change pH

Buffers are created by placing large amounts of a weak acid/base into a solution with its conjugate (salt)

Weak acid and conjugate base remain in solution together without neutralizing each other

Presence of the conjugate pair makes the buffer effective

If enough strong acid/base is added that all of the acid or conjugate base is reacted, the buffer breaks

Higher concentrations of the conjugate pair resist pH change (better buffers) better than lower concentrations

Henderson-Hasselbalch

When concentrations of acid and conjugate base in a solution are the same, pH=pKa and pOH=pKb

Choosing an acid for a buffer solution requires choosing an acid with a pKa close to the desired pH

(almost equal amounts of acid and conjugate base; makes buffer flexible in neutralizing both added H+ and OH-)

Buffers cannot be created from a very strong acid and its conjugate base because the conjugate base will be very weak and will not readily accept protons

Indicators

Weak acids which change colors in certain pH ranges due to LeChatelier’s principle

HIn ⇄ H+ + In-

Ka = [H+][In-]/[HIn]

Protonated HIn state must be a different color from deprotonated In- state

Acidic environment causes excess H+ to drive equilibrium to the left (color of HIn); basic environment causes excess OH- to react with H+ from indicator and drive reaction right (color of In-)

Color change occurs when [HIn] = [In-]; or pH = pKa

Choose an indicator whose pKa matches the pH at the titration’s equivalence point

Titration

Neutralization reactions are performed by titration, where a base of known concentration is slowly added to an acid or vice versa

Titration curves

Midpoint also called half equivalence point occurs when [HA] = [A-] (pH = pKa)

Equivalence point occurs when just enough base has been added to neutralize all the acid initially present (equimolar)

HA, A- present before midpoint; A- at midpoint, OH- after midpoint

X. Laboratory Overview

Weighing hot objects on a scale creates convection currents, making object appear lighter than it truly is

Not rinsing a buret in a titration leads to it being diluted

[Image: Molecular geometry (VSEPR) diagram showing various shapes such as linear, trigonal planar, tetrahedral, etc.]

III. Intermolecular Forces and Properties (continued)

Polynomial (assorted) details shown in figure

Relationship of geometry to polarity and IMF

Common Polyatomic Ions
C2H3O2- Acetate
NH4+ Ammonium
CO3^2- Carbonate
ClO3- Chlorate
ClO4- Perchlorate
CrO4^2- Chromate
CN- Cyanide
Cr2O7^2- Dichromate
HCO3- Bicarbonate
HSO4- Bisulfate
HSO3- Bisulfite
OH- Hydroxide
ClO- Hypochlorite
NO3- Nitrate
NO2- Nitrite
C2O4^2- Oxalate
MnO4- Permanganate
PO4^3- Phosphate
SO4^2- Sulfate
SO3^2- Sulfite

IV. Chemical Reactions

Types of reactions

Synthesis: everything combines to form one compound

Decomposition: one compound + heat is split into multiple elements/compounds

Acid-base rxn: Acid + base -> water + salt

Oxidation-reduction (redox) rxn: changes the oxidation state of some species

Combustion: substance w/ H and C + O2 -> CO2 + H2O

Precipitation: aqueous solutions -> insoluble salt (+ more aq sometimes)

Can be written as net ionic - Those free ions not in net ionic are spectator ions

Solubility rules

Alkali metal cations or ammonium (NH4+) cations are ALWAYS soluble

Compounds with a nitrate (NO3-) anion are ALWAYS soluble

Common polyatomic ions

Calculations

Percent error: 100 * abs(experimental - expected)/(expected)

Combustion analysis - use law of conservation of mass

Gravimetric analysis - determine identity by comparing mass percentages

Oxidation states

Neutral atoms 0

Monoatomic ions equal to ion charge

Oxygen -2 (except in H2O2 where it is -1)

Hydrogen +1 with nonmetals; -1 with metals

In absence of oxygen, most electronegative element takes usual charge

C, N, S, P vary often

Redox reactions

Write 2 half-reactions; balance O with H2O; balance H with H+; combine

ACIDIC: stop here

BASIC: add OH- to both sides to neutralize H+; cancel H2O

Acids and bases (briefly)

Color change signals end of titration

Acids donate H+, bases donate electrons

Water is amphoteric

V. Kinetics

Rate law: Rate = k [A]^x [B]^y [C]^z

Determine x, y, z via experiments

Keq = forward/backward rate constants

Units: rate M/s; k units depend on order

Reaction stoichiometry relation for rates

Orders: zero, first, second order rules; half-life

Collision theory

Activation energy Ea; orientation requirements; temperature and concentration effects

Reaction energy profile

Reaction mechanisms

Catalysts

VI. Thermodynamics

Temperature/heat

Temperature = average kinetic energy

Heat = energy flow between systems

First law

Bond formation/releases energy

Exothermic vs Endothermic

Energy diagrams

Enthalpy

Formation; Hess’s law; bond energy; enthalpy of solution

Hydration energy; Coulombic energy

Phase changes

Sublimation/Deposition; boiling point; fusion/vaporization; IMF role

Calorimetry

Specific heat; q = mcΔT; q1 = q2; calorimetry for reactions; delta H

Heating curves

VII. Equilibrium

Keq; rate concepts; Le Chatelier; Q vs Keq

K values: Kc, Kp, Ksp, Ka, Kb, Kw

Manipulating Keq

Le Chatelier's principle specifics

Solubility and common ion effects

VIII. Acids and Bases

pH and related formulas

Strong vs weak acids; acid strength; Ka, Kb; RICE table

Kw relations; pH/pKw; buffer concepts; Henderson-Hasselbalch

Buffers and indicators; titration concepts; equivalence and midpoint

IX. Applications of Thermodynamics

Entropy; standard entropies; delta S; Gibbs free energy; Delta G; spontaneity

Galvanic/voltaic cells; anode/cathode; electron flow; salt bridge

Nernst equation; non-standard conditions; cell potential

Electrolytic cells; calculation of metal deposition

Voltage and favorability

X. Laboratory Overview

Practical notes: weighing hot objects; rinsing burets

The content above mirrors typical AP Biology topics with detailed subpoints and examples as presented in the source document, including electron configuration, bonding types, thermodynamics, equilibria, acids/bases, and laboratory principles. All sections, subsections, and data points are preserved to enable reconstruction of the original document from this JSON export.

Image reference: Molecular geometry (VSEPR) diagram showing various shapes like linear, trigonal planar, tetrahedral, trigonal pyramidal, square planar, etc. {#figure-molecular-geometry}

Alt text: “Molecular geometry (VSEPR) diagram illustrating common geometries and steric numbers”

Caption: “Figure: Molecular geometry examples according to VSEPR theory.”

[Image] url: images/molecular-geometry-vsepr-diagram.jpg; alt: Molecular geometry diagram per VSEPR; caption: Figure: Molecular geometry (VSEPR) illustration

AP Biology Review: Atomic Structure to Thermodynamics

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Summary & Guide