Entire Chem 1C Lecture/Textbook Notes.doc

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University of California - Irvine
Amanda Brindley

Chem 1 C 1 Chap 14 Chemical Equilibrium • • dynamic equilibrium • o the state of a reaction in which its forward and reverse reactions occur at equal rates so that the concentration of the reactants and products does not change with time. • chemical equilibrium • o rate of forward and reverse reactions are equal and the concentrations of reactants and products remain constant • physical equlibrium • o equilibrium between two phases of the same substance o  ex. ice with water • equilibrium constant • o unitless, its an activity • o o  equilibrium constant = (concentration of products (M)) / (concentration of reactants (M))  instead of concentration, partial pressures can be used. that would be Kp (if gas) and instead of Molar, Pressure is used  treat liquids and solids as constants; they are not part of the concentration   only gas and aqueous contributes to the equilibrium constant o o  partial pressure equilibrium constant = concentration equilibrium constant (.0821T)^(moles of gaseous products - moles of gaseous reactants)  deriving this equation:  Chem 1 C 2  o Combining K for multiple reactions o  similar to hess's law, but not the same  • Le Chatelier's principle • o if an external stress is applied to a system at equilibrium, the system adjusts in such a way that the stress is partiall offset as the system reachers a new equilibrium position o  concentration changes   does not change equilibrium constant, only position  add more products (right) will cause the equilibrium to shift to the left (reactants) because more reactants can be made  volume/pressure changes   does not change equilibrium constant, only position  pressure increase (volume decrease) favors will shift toward the side w/a smaller # of net gaseous moles Chem 1 C 3  if both numbers r the same (i.e. one mole vs one mole) changing pressure/volume does nothing to the equilibrium position  adding pressure w/other gasses does nothing  temperature changes   changes equilibrium constant  adding heat to an endothermic (delta H is +) process (that is, in one direction) will cause the equilibrium constant to rise or fall depending on the coefficient of the products and reactants and its place in the concentration constant  ex: heat added to    AB --> 2CD is endothermic (detla H is +, with AB), so adding heat to AB will make more CD and therefore the K will increase  adding a catalyst   does not change equilibrium, just the rate it reaches it • Free energy and Equilibrium • o o  change in free energy = -(8.31J/molK)(Temp)(lnK) o o  change in free energy = change in enthalpy - (temperature)(change in entropy) o those two equations combined makes von't Hoff exn o Chem 1 C 4 o Chem 1 C 5 Chem 1 C 6 Chap 15 Acids and Bases • pH • o -log of the hydrogen ion concentration o  pH = -log[H3O+] = -log[H+]  pOH = -log[OH-]  pH + pOH = 14   o Acid o  pH = 6.99-  [H+] > 10^-7  denotes a photon or electron pair  has a conjugate base    H3O+ is the strongest acid that can exist in an aqueous solution o Base o  pH = 7.01+  [H+] < 10^-7  accepts a photon or electron pair  has a conjugate acid    OH- is the strongest base that can exist in an aqueous solution o neutral o  7.00  [H+] = [OH-]= 10^-7  Kw = [H+][OH-] = [H3O+][OH-] = 10^-14 • Strength of acids and bases • o strong Acids and Bases o Chem 1 C 7  electrolytes that are assumed to ionize completely in water    most acids and bases are weak   ionize only to a limited extent in water  ICE chart   if a weak acid or base has a small [H] or [OH] concentration (smaller then 10^-6), then the initial concentration of [H] or [OH] needs to be used (10^-7). otherwise, it can be assumed that this initial concentration does not affect outcome. • ionization constant • o Ka - acid o Kb - base o only applies to weak acids/bases o  because the reactants of a strong acid/base ionize completely, it would be like dividing by 0 o water can be ignored most of the time because it has such a low H+ ionization o o  relationship between Kw (10^-14) and the Ka and the Kb of a solution • percent ionization • o o is used to check weak acid/base concentration (%ionization is < 5%, the approximation is valid) o if %ionization ~ 10%, only 1 out of every 9 acid/base molecules ionize • diprotic and polyprotic acids • o with an acid such as H3PO4, the Ka concentration can be assumed to not hav HPO4 or PO4 Chem 1 C 8 o  H3PO4 -> H + H2PO4  HPO4 and PO4 are made, just in much smaller amounts • Molecular Structure and Strengths of Acids • o stronger the bond, weaker the acid o  if two acid molecules that do not want to break apart creates a weaker acid o  as electronegativity goes up, so does bond strength which means weaker acid o   increasing acid strength, decreasing electronegativity (hydrohalic acids)  HF < HCl < HBr < HI o higher polarity, stronger acid o  for example, if one end of a molecule is very negative while the other is very positive, they will be more likely to break apart o Oxoacids (central element surround by OH's) o  electronegativity rule is opposite  as electronegativity of central atom goes up, OH's bond strength goes down, resulting in a stronger acid   HBrO3 < HClO3  acid strength increases as oxidation number of central atom increases   HClO2 < HClO3 < HClO4 o Carboxylic Acid o  more stable anion (conjugate base), stronger acid  high resonance, aromaticity, higher electronegative = stabler atom = stonger acid   H2Se << H2S < H2O  H2SeO4 < H2SO4  CH2ClOOH < CHCl2COOH • Properties of Acid/Base Salts • o salt o  compound formed by the reaction of an acid and base  strong electrolytes that completely dissociate into ions in water, which is called salt hydrolysis Chem 1 C 9  usually affects pH of a solution o neither salt accepts or donates an H+ o  produce neutral solutions  NaNO3 -> Na + NO3   NO3 will not accept an H because it is a strong acid o strong base and a weak acid o  produces basic solutions  CH3COONa -> Na + CH3COO   strong base will steal H from water leaving OH- o strong acid and weak base o  produce acidic solutions   NH4Cl -> NH4 + Cl   strong acid Cl will have no affinity for H, but NH4 will donate H to become its weak conjugate acid NH3 o both acid and base are weak o  Kb > Ka - basic   if the Kb for the anion is greater then Ka for the cation, the solution will be basic because the anion will hydrolyze more and there will be more OH  Ka > Kb - acidic   if the Kb for anion is smaller the Ka for the cation, the solution will be acidic because the cation will hydrolyze more and there will be more H  Kb = Ka - neutral • acid/base definitions • o Arrhenius - acid donates H, base donates OH o Bronstead - acid donates H, base accepts H o Lewis - acid accepts lone pair, base donates lone pair o  Ag + (acid) + 2NH3 (base) -> Ag(NH3)2 +  Cd 2+ (acid) + 4I - (base) -> CDI4 2+  Ni (acid) + 4CO (base) -> Ni(CO)4 o Henderson-Hasselbalch equation (only works for buffer problems) Chem 1 C 10 o o  pH = -logKa + log[conjugate base]/[acid] • Buffer solution • o when a weak acid or base is present with its salt, the solution has the ability to resist changes in pH upon the addition of small amounts of either acid or base • identifying if something is an acidic or a basic • o acidic o  if the cation+ is derived from a weak base o basic o  if the anion- is derived from a weak acid o no effect o  if either are derived from a strong acid or base Chem 1 C 11 Chem 1 C 12 Chap 16 Acid-Base Equilibria and Solubility Equilibra • common ion effect - shift in equilibrium caused by the addition of a compound having an ion in common with the dissolved substance • o ex: o  CH3COONa -> CH3COO- + NA+ strong, dissociates completely in H2O  CH3COOH CH3COO- + H weak acid  if we add these two reactions together, the amount of CH3COO- from CH3COONa will suppress the ionization of CH3COOH (that is, will shift the reaction left because CH3COO- was added to the right side)  this will make the concentration of H less and result in a less acidic solution  CH3COO- is the common ion o pH = -logKa + log[conjugate base]/[acid] • Buffer Solution • o when a strong acid/base is added to a weak base/acid, the weak base/acid will create salt (conjugate acid/base) o  if its a weak acid/base with a weak base/acid, use common ion effect formula (since its the same thing) o when a weak acid or base is present with its salt (conjugate base/acid), the solution has the ability to resist changes in pH upon the addition of small amounts of either acid or base because: o  a strong base can't exist in a solution w/weak acid (visa versa) o  o valid way to make NH3/NH4 buffer o  add equal amounts (volume and M) of NH3 and NH4 together  mix 1 M NH3 with half as much 1 M HCl   half of NH3 will be converted to HCl o examples of buffers: o Chem 1 C 13   does not include strong acid/bases   KCl/HCl. o we can make a solution a certain pH o  [conjugate base]/[acid] = 1 (or visa versa)   pH = pKa  chooses a weak acid or base which has a pKa close to desired pH  ex 16.4 in book • acid-base titrations • o types: o  strong A/B or B/A    no buffer area  at S, [H] = [OH] = 10^-7  strong B/weak A, strong A/weak B  Chem 1 C 14   buffer area  at S, all weak base has been converted to weak acid  weak A/B or B/A (not gunna do these because they are confusing) o S is equivalence point, not always at pH = 7 o  moles acid = moles of base  moles of OH added = moles of H orginally present o How to solve titration problems, depending on stage o   1: weak acid or base problem   (in this picture it would b a weak base problem)  2: buffering problem   strong acid/bases do not have buffering stage  pH = pKa + log(base/acid)  3: equivalence point problem Chem 1 C 15   moles acid = moles of base  moles of OH added = moles of H originally present  all weak base has been converted to a weak acid (or visa versa)  treat as a weak acid problem, use Ka and adjust concentrations (m1v1=m2v2)  4: strong acid or base problem   (the picture yields a strong acid problem)  make sure to subtract what was used to convert the weak base to the weak acid (or visa versa) when finding the concentration • acid-base indicators • o has distinctly different colors in nonionized and ionize form o end point - occurs when the indicator changes color o an indicators equation and why it changes color o  HIn(aq) <=> H+(aq) + In-(aq)  where HIn and In- have different colors  if the indicator is in a sufficiently acidic medium, more HIn will be present  basic solution will cause more In- and that ion's color will show • Solubility • o Ksp - solubility constant o just like equilibrium in that o  products/reactants  solids aren't included  concentrations (of ions and cations = molar solubility) are raised to the power of their coefficients o saturated solutions o  solubility (g/L) --> molar solubility (mol/L) --> Ksp of compound o predicting precipitate o  find Q, just like in Keq, to find which way the reaction will go   no precipitate   Q
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