AP Chemistry Review: Notes and Study Guide PDF

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 weighte...

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 - necessary 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  ​ Vapor is collected, cooled, and condensed back to a liquid separate of leftover liquid ​ 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 (different gases will have same KE at same temp) ○​ 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 (gas molecules are packed too tightly together) ​ Volume of gas molecules becomes significant (less free space for molecules to move around than predicted) ​ Gas molecules attract one another and stick together (real pressure is smaller than predicted pressure) ○​ 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 2.998*10^8; v = frequency; lambda = wavelength  ○​ Beer’s law: A = abc ​ A = absorbance; a = molar absorptivity (constant depending on solution); b = path length of light through solution (constant); c = concentration of solution ​ Colorimetry - direct relationship b/w 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 ​ Species with the H+ ion are acids, same species but without H+ is a base - conjugate acid/base pairs ○​ 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 (more likely to collide) ○​ Stirring increases reaction rate for heterogeneous mixtures (causing heterogeneous mixture to move around increases collisions; insignificant once the mixture becomes homogeneous due to the number of collisions happening due to inherent motion of aq molecules)  ○​ Greater temp increases reaction rate (greater fraction of reactant molecules has sufficient energy to exceed activation energy barrier - vertical line on Maxwell-Boltzmann w/ multiple temps) ○​ Reactions only occur if reactants collide with correct orientation to break the right bonds ​ Reaction energy profile ​ Reaction mechanisms ○​ Species that are produced in a mechanism but are also fully consumed and do not appear in the balanced equation are intermediates ○​ Adding up all mechanism steps and canceling out different species leads to the balanced rxn ○​ Elementary steps w/ 2 reactants (even if they are the same) are bimolecular; elementary steps 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 ​ Make rate for slow step (ex. If X + B -> Y, rate = k[X][B]) ​ Substitute in rate for X from above equilibrium reaction ​ Compare to actual reaction’s rate equation ○​ Slow step has highest activation energy ​ Catalysts ○​ Catalysts increase rate of chemical reaction without being consumed in the process ○​ Catalysts do not appear in balanced equation ○​ In a reaction mechanism, catalysts enter first, then exit  ○​ Catalysis (reaction with a catalyst) ​ Surface catalysis - reaction intermediate is formed ​ Enzyme catalysis - catalyst binds to reactants to reduce activation energy ​ Acid-base catalysis - reactants lose/gain protons to change reaction rate VI.​ Thermodynamics ​ Temperature/heat ○​ Temperature is the average amount of kinetic energy due to molecular motion in a given substance ○​ Heat is the energy flow between 2 different substances at different temperatures ○​ First law of thermodynamics: energy can be neither created nor destroyed ○​ When bonds are formed, energy is released; when bonds are broken, energy is absorbed ○​ Exothermic - energy transferred from system to surroundings (delta H is negative) ​ More energy is released when the product bonds form than is necessary to break reactant bonds ○​ Endothermic - energy transferred from surroundings into system (delta H is positive) ​ More energy is required to break reactant bonds than is released when bonds in products form ​ 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 ​ Multiply delta Hf for each product/reactant by the coefficient  ​ 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 since bonds are being broken) ​ 2: Separation of solvent molecules - water molecules must spread out to make room for the solute ions (requires energy to weaken the IMF between water molecules - positive delta H) ​ 3: Creation new attractions - free floating ions are attracted to the dipoles of water molecules (energy is released - 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 - lower atmospheric pressure (high elevation) means a lower boiling point ○​ 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 for a solid than a liquid, 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 ○​ Heat added (J or cal) 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 ​ In reaction 2A -> B, concentration of A will decrease 2x as much as concentration of B increases ○​ Equilibrium expression/law of mass action ​ 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 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 ration 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 (-> not ⇄) so 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 (hydrogen ion) from an acid (protonates/ionizes, not dissociates) ○​ 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 easier it is for H+ ions to break free), the stronger the acid is ​ Lower concentration -> higher percent dissociation; a greater 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 a (usually O) atom at the end of the molecule, making it easier for it to detach ​ H in a hydroxyl group (-OH) are dissociable due to O being much more electronegative than H ​ H bonded to C is almost never dissociable since H and C have similar electronegativity values and share their electrons equally ○​ 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 = 1*10^-14 = Ka*Kb ​ 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 > 1, Q = 1 ○​ Non standard conditions ​ If Q=Keq, cell voltage drops to 0; increasing Q decreases cell potential and vice versa ​ Overall potential decreases as a reaction progresses (product conc increases, reactant conc decreases) ​ Nernst equation: Ecell = E0cell - (RT/nF)*ln(Q) where n is # of electrons transferred (always positive) ​ Electrolytic cells ○​ Outside source of voltage is used to force an unfavored redox reaction to occur ○​ Occur primarily in aqueous solutions (chemical dissolved in water; ion/water molecule is oxidized/reduced) ○​ Compare reduction potential of cation with that of water (reduction) to determine which is reduced; compare oxidation potential of anion with that of water (oxidation) to determine which is oxidized ​ Then balance the 2 oxidation/reduction half reactions to form one net ionic equation ​ value for cell potential from the half reactions should always be negative ○​ Oxidation at + anode (left); reduction at - cathode (right) ​ Signs change from galvanic cell setup ○​ Electrons flow from anode to cathode (L to R) ​ + to - (instead of - to + like in galvanic cells) ○​ Used for electroplating ​ I = q/t ​ Moles of electrons = coulombs/ 96500 C/mol ​ Moles of metal from moles of electrons (from metal half reaction) ​ Moles of metal -> grams ​ Voltage and favorability ○​ Redox reaction is favored if potential is positive ○​ Delta G = -n*F*E0 (n is positive # of electrons transferred, F is 96500, E0 is standard cell potential in V = J/C) 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