Chem 161

Acids-base descriptions…

• Arrhenius description: H+ (acid) and OH- (base)

• Bronsted description: Proton donors (acid) and proton acceptors (base)

• Lewis description: electron acceptors (acid) and electron donors (base)

Acid strength…

• Adding oxo groups increases acid strength

  • EN O pulls electron density away from central atom making it more positive and easier to donate proton (H+)

  • More oxo groups means more resonance structures, making c.b more stable and since neg. Charge from loss of proton after donation can be delocalized over multiple oxygen hence less likely to revert back to protonated state (stable c.b. = stronger acid)

  • Note that more stable the c.b, the weaker the c.b. Strong c.b. Wants to gain H+ so should be unstable. 

Acids in disguise…

• Metal(H2O)₆ⁿ⁺: metal complex ions, central metal surrounded by water, looks neutral but acts as acids by donating H+ from H2O molecules

• Metal ion polarized O-H bonds of H2O, making it easier for them to lose H+

• Metals with higher change + smaller size have stronger acidity because they exert stronger pull on H2O electrons

Acid and Base Properties of Salts:

K_w = K_a x K_b

pK_a = -log (K_a) / K_a = 10^-pKa Polyprotic acids: 

Polyprotic acids can supply for than one proton (has multiple H+) so they dissociate in stepwise manner, one proton at a time

• Ka1 > Ka2 > Ka3 > etc

• Exception: H2SO4^- -> H+ + HSO4- ⇄ H+ + SO4^2-

Ka, Kb, and Kw:

Stronger acids -> higher Ka

Stronger bases -> higher Kb

Autoionization of water

• K_w = K_a x K_b

• K_w = 1.0 x 10^-14

Calculating the pH of weak acid mixtures: 

Common ion effect is shift in equilibrium position that occurs because of addition of ion already involved in equilibrium reaction (le chatelier principle)

Titrations of Strong Acids and Bases:

M₁V₁ = M₂V₂

Dilution does not affect number of moles of solute BUT reaction can change number of moles

Final moles H+ = Initial moles H+ - Initial moles OH-

Remember to consider FINAL volume of reaction

Equivalence point: when amount of base (titrant) added is equal to amount of analyte (acid) in sample

• Basically when all of acid reacts with base, so moles of acid and base reaction in stoichiometrically equivalent amounts

• No excess titrant left in solution

• For strong acids and bases, pH is 7 (neutral) 

Titrations of Weak Acids with Strong Bases:

Find pH via stoichiometry calculation (amount of OH- added reacts completely) then equilibrium calculation (remaining amount of acid is used for this)

Equivalence points of strong acids produce neutral salt so is 7

Equivalence points of weak acids have pH greater than 7 because resulting solutions contains basic salt

Always think about what is IN THE SYSTEM, including PRODUCTS

Introduction to Buffers:

Buffers contains a weak acid and it c.b OR weak base and its c.a

• Buffer region is where system is resistant to change in pH

Why do Buffers Buffer?:

Henderson-Hasselbach: pH = pKa - log ([HA]/[A-]) which is acid/base OR pH = pKa = log ([A-]/[HA]) which is base/acid

Note that buffers react with incoming H+ or OH- and stabilize pH of system

The Anatomy of a Buffer:

Buffers are pH dependent, and pH is determined by ratio of acid to base

Amount of H+ and OH- buffer can absorb based on concentration (low concentration = small capacity and larger pH changes, high concentration = large capacity and smaller pH changes)

• Because higher the concentration of acid/base in bugger, the greater its ability to neutralize hence smaller pH change

pKa is pH at which weak acid/base is 50% dissociated so concentrations of acid and c.b (or vice versa) is equal

• At pH = pKa, bugger is most effective because concentrations of acid/c.b are equal, allowing buffer to neutralize equal amounts of added acid/base