ap chem acids and bases

strong acids

HCl (hydrochloric)

HBr (hydrobromic)

HI (hydroiodic)

HNO₃ (nitric)

HClO₄ (perchloric)

H₂SO₄ (sulfuric)

weak acids

anything else with a H in the front

strong bases

hydroxides of Group I and IIA

ex. NaOH, Ca(OH)₂

weak bases

NH₃ (ammonia)

CH₃NH₂ (methylamine)

C₂H₅NH₂ (ethylamine)

neutral salts

salts that are formed from the cation of a strong base and the anion from a strong acid

ex. NaCl, KNO₃, LiCl

acidic salts

salts that are formed from the cation of a weak base and the anion from a strong acid

ex. NH₄Cl, NH₄NO₃

basic salts

salts that are formed from the cation of a strong base and the anion from a weak acid

ex. NaH₂PO₄, NaC₂H₃O₂

the weaker the acid

the stronger its conjugate base

vice versa, strong acids have very weak conjugate bases

• Weak acids have strong conjugate bases (will produce more OH- ions than just water) because the acid form of the molecule is more stable than the anion form.

  
  

the larger the K_a or K_b

the stronger the acid/base

BUT strong acids do not usually have K_a bc [HA] is so small alr

monoprotic diprotic polyprotic

mono: can only donate one proton (HF)

diprotic: yields two H⁺ per molecule but only ionizes one at a time (H₂CO₃)

polyprotic: capable of losing more than a single proton per molecule (H₃PO₄)

• will only lose ONE hydrogen at a time. It will take more energy to lose (donate) the next hydrogen, so therefore it will be a weaker acid, (smaller K_a)

oxyacids

most acids are oxyacids, acidic proton is attached to O

acid strength inc with increasing numbers of oxygen atoms

• electronegative O draws electrons away from O-H bond, O-H bond becomes weaker and H⁺ easier to get released

oxyacids electronegativity

acid strength increases with increasing electronegativity of oxyacids

• HOCl > HOBr > HOI > HOCH₃

hydrohalic acids

HCl, HI, HF

• HF is only weak hydrohalic acid. Although the H-F bond is very polar, the bond is so strong (due to the small F atom) that the acid can not completely dissociate

amphoteric substance

substance that can act as an acid or as a base

ex. H₂O, NH₃, HSO₄^-

autoionization of water

K_w = [H₃O⁺][OH⁻]
                                                               
At 25^oC, K_w = 1 × 10⁻¹⁴ because [H₃O⁺] = [OH⁻] = 1 × 10⁻⁷ M

pH = pOH = 7

K_w only changes with

temperature, always 1 × 10^-14 at 25 degrees

relating pH with endothermic/exothermic

ex. If the pH of a water at 50 degress is 6.7, the ionization of water is endothermic, because pH < 7 indicates more H⁺ ions are produced (forward rxn with prod. products favored). If adding heat favors the products, heat acts like a reactant, so the ionization of water is endothermic.

exceptions to calculating pH of strong acid solutions!

1. If acid concentration is less than 1.0 × 10^-7, water becomes the important source of [H⁺] and the pH is 7

• pH of acidic solution can’t be greater than 7 at 25 dg

• say “since the [H⁺] is less than 1.0×10^-7, the [H⁺] from the acid is negligible and the pH = 7.00

2.   Another exception is calculating the pH of an H₂SO₄  solution that is more dilute than 1.0 M.  At this concentration, the [H₃O⁺] of the HSO₄⁻ must also be calculated.

pH of weak acids

will need equilibrium, x²/(i - x)

• if you know x, use it dont neglect, (ex. given pH)

acetic acid AKAs

HC₂H₃O₂ or CH₃COOH, aka ethanoic acid

determining the pH of a mixture of weak acids

Only the acid with the largest K_a value will contribute an appreciable [H⁺].

• if strong + weak only care about strong

For a weak acid/base, percent dissociation (or ionization)

Increases as the acid/base becomes more dilute. Equilibrium shifts to the right

+ vice versa

acidic salts reasoning

ex. NH₄Cl

NH₄⁺ is conjugate acid of a weak base, it will hydrolyze in water producing H⁺ ions, pH below 7

NH₄⁺ + H₂O ↔ NH₃ + H₃O⁺

basic salts reasoning

ex. NaF

F^- is conjugate base of a weak acid, it will hydrolyze in water producing OH^- ions, pH above 7

F^- + H₂O ↔ HF + OH^-