Similar document for AP Bio: AP Bio Review
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 ewill 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 eare 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 efor 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 eto 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 eexperience greater LDF (more random motion) ■ Higher molar mass usually means greater LDF (as mass increases, eincreases 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 6.62610^-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 (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 OHto 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<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 = 110^-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 (moles of gas) ■ Increasing volume decreases pressure and vice versa ■ Adding a non-reacting gas (noble 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 Q
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 OHand 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-] = 1.010^-14 at 25 C for any aqueous solution ■ pH + pOH = 14 ■ Kw = 110^-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 <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 Instate ○ Acidic environment causes excess H+ to drive equilibrium to the left (color of HIn); basic environment causes excess OHto 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, Apresent before midpoint; Aat midpoint, OHafter midpoint
IX. Applications of Thermodynamics ● Entropy ○ Measure of randomness or dispersion of the system ○ 0 entropy is a solid crystal at 0K (has never been reached experimentally) ○ Standard entropies are calculated at 25 C ○ Standard entropy change delta S = S for products S for reactants ○ If left side of a reaction has more motion, delta S is negative; if right side has more motion delta S is positive ● Gibbs free energy
○ Delta G = Gf of products Gf of reactants ○ Negative delta G is spontaneous (thermodynamically favored); positive delta G is nonspontaneous (thermodynamically unfavored); delta G = 0 means rxn is at equilibrium ○ Delta G = delta H T*delta S (T in K) ○ Favorability
○ Delta G in phase changes is 0 since at a normal phase transition temp, the substance is equally stable in either phase ■ Boiling/melting point can be solved for knowing delta H and S since delta G = 0 ○ Delta G = -RTlnK (R = 8.314, T in K, Keq) ○ The larger the reduction potential on a half reaction, the more likely it is to occur ● Galvanic/voltaic cells ○ Favored redox rxn generates a flow of current ○ Oxidation at anode (left); reduction at + cathode (right) ○ Electrons flow from anode to cathode (L to R) ■ Electrons released from oxidation pass to chamber to be consumed in reduction ○ Flow of electrons creates current ○ Salt bridge between two cells maintains electrical neutrality ■ positive cations from salt bridge solution flow to cathode which is losing positive charge (needs + to balance); negative anions from salt bridge solution flow to anode which is gaining positive charge (needs to balance)
○ Under standard conditions (1M solutions, 1 atm, 25 C), cell voltage is the same as total redox voltage ■ Keq >> 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 = -nF*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