pH

• Concentration equations are usually in pH

  • pH = -log[H+]

  • pOH = -log[OH-]

  • pKa = -log[Ka]

• In a solution

  • If [H+] = [OH-] the solution is neutral with a pH is 7

  • If [H+] > [OH-] the solution is acidic wih a pH of <7

  • If [H+] < [OH-] the solution is basic with a pH >7

• Increasing pH means decreasing [H+] and making a solution less acidic.

• Decreasing pH means increasing [H+] and making a solution more acidic.

• Some hydroxide salts will not dissolve in high pH because of the common ion effect.

  • In the equation Mg(OH)2 (s) ⇌ Mg2+ (aq) +2OH-(aq)

  • If the salt is added to a solution with significant hydroxide ions in it, it will shift to the left and decrease the amount of salt dissolving. If added to a solution with an abundance of hydrogen ions, it will dissolve more because the hydrogen and hydroxide will react, creating less hydroxides in the solution, and the salt to dissolve more.

Acid Strengths

Strong Acids

• Strong acids will dissolve completely in water and never reach equilibrium. There is no equilibrium constant for strong acids or bases because of no equilibrium

  • Important strong acids

    • HCl, HBr, HI, HNO3, HClO4, H2SO4

  • Important strong bases

    • LiOH, NaOH, KOH, Ba(OH)2, Sr(OH)2

• Because completion occurs, the conjugate base of the acid must be extremely weak.

• Finding the pH of strong acids is easier than finding the pH of weak acids because strong acids will dissociate completely while weak acids will only partly dissociate so the concentration of hydrogen ions is different in the beginning versus the end

Weak Acids

• A weak acid in water will have a small fraction dissociate into the conjugate base and hydrogen ions while the rest remains undissociated aqueous particles.

• Weak acid and base dissociation constants are Ka and Kb, specific to only acids and bases

• Ka = [H+][A-]/[HA]

  • [H+] is hydrogen ion concentration in molarity

  • [A-] is conjugate base ion concentration in molarity

  • [HA] is concentration of undissociated acid molarity

• Kb = [HB+][OH-]/[B]

  • [HB+] is conjugate acid ion concentration

  • [OH-] is hydroxide ion concentration

  • [B] is unprotonated base molecules

• Greater Ka value means more acid dissociation (stronger acid)

• Greater Kb value means more protonatization (stronger base)

  • Bases do not dissociate but accepts protons

Percent Dissociation

• More hydrogen ions an acid can donate means stronger acid and the ease of which an acid can dissociated is determined by molecular structure

• For binary halogen acids, HI, HBr, and HCl are all strong acids but HF is not because of its high electronegativity

• For oxacids, the reverse is true

  • For HOF, the hydrogen and fluorine are on opposite ends, so the hydrogen is less electron dense because the oxygen acts as a buffer and can dissociate

  • For HOBr, the hydrogen is slightly more electron dense than the hydrogen in the HOF because Bromine isn’t as electronegative as the fluorine

• The more oxygens you add to an oxacid, the stronger of an acid it becomes.

• If there is an abundance of water with an acid, or a smaller concentration of acids, the acid can find a water to donate the hydrogen to and will react/dissociate more

• If there is a greater concentration of an acid, or less water molecules, it will not be able to dissociate as much because there is less water to be able to accept a hydrogen ion

Acid/Base Structure

• Only some hydrogens in an acid are able to dissociate

  • This is why acetic acid is written as HC2H3O2 rather than C2H4O2

• The hydrogen will dissociate from the least electronegative end.

• However hydrogens attached to a carbon nearly never dissociate because they share the electron fairly equally because of the similar electronegativity.

• Generally, hydrogens that are written in the beginning of an acid are able to dissociate while ones in the middle cannot

pH and Solubility

• Mg(OH)2 (s) ⇌ Mg2+ (aq) +2OH-(aq)

• If the salt is placed in a high pH, it would be less soluble than if it was placed in a lower pH.

  • Higher pH means more hydroxide ions and cause the salt to be less soluble/react less

  • Lower pH means more hydrogen ions and cause the salt to dissociate more/be more soluble

• This is similar to the common ion effect

Polyprotic Acids

• Polyprotic acids can give up more than one hydrogen in a solution

  • H2SO4 and H3PO4 are polyprotic

• The first proton is easier to give up than the second one

  • The first Ka for H3PO4 is 7.1x10^-3

  • The second Ka for H2PO4 is 6.3x10^-8

  • The third Ka for HPO4 is 4.5x10^-13

• H3PO4 is stronger of an acid thand H2PO4 because the protons are more attracted to the H2PO4 and not as easily given up

Equilibirium Constant of Water (Kw)

• Water comes to equilibrium through the reaction

  • H2O (l) ⇌ H+(aq) + OH-(aq)

    • Kw = 1x10^-14 at 25 degrees celsius

  • Kw = 1x10^-14 = [H+][OH-]

  • pH + pOH =14

• The H and OH in any acid or base solution must be consistent with the ionization of water

• Ka and Kb values are also consistent with the water equilibrium reaction

  • Kw = 1x10^-14 = Ka x Kb

  • pKa + pKb =14

• pHs can exceed 14 and be lower than 0, but are uncommon

• The Kw constant is only at 1x10^-14 at 25 degrees celsius because bonds will form and break as the temperature changes so the Kw value changes

Neutralization Reaction

• Neutralization reactions occur when an acid donates its proton to a base

  • Strong acid + strong base

    • Both substances dissociate completely and the only important ions in this reaction are the hydrogen and hydroxide

    • Ex: HCl + NaOH

    • This always results in the creation of water and spectator ions

  • Strong acid + strong base

    • The strong acid dissociates completely and donates a proton to the weak base with the product being the conjugate acid

    • HCl +NH3

  • Weak acid + strong base

    • Strong base accepts protons from weak acid. Creates conjugate base of weak acid and water

    • HC2H3O2 + NaOH

  • Weak acid + weak base

    • Simple proton transfer

    • HC2H3O3 + NH3

Buffers

• A buffer is a solution with a stable pH. Adding an acid, base, or water to a buffer solution will not greatly affect the pH.

• A buffer is created by adding a significant amount of weak acid or base along with it’s conjugate as salt. This combination will not neutralize

  • NaC2H3O2 and HC2H3O2 added together will create a buffer solution.

  • If strong acid was added to this buffer solution it would create more acetic acid

  • If a strong base was added it would create more of the acetate ion

• More concentration of the weak acid/base and its conjugate will create a stronger buffer solution

• Henderson-Hasselbach equation can calculate exact pH of a buffer

  • pH = pKa + log([A-]/[HA])

    • [HA] = molar concentration of undissociated weak acid

    • [A-] is molar concentration of conjugate base

  • pOH = pKb + log([HB+]/[B])

    • [B] is molar concentration of weak base

    • [HB] is molar concentration of weak acid

• When choosing an acid for a buffer solution, pick one with a pKa that is close to the desired pH.

• A buffer cannot be created through a strong acid and its conjugate because it would dissociate completely. Same is said for bases.

Indicators

• Indicators are weak acids that change colors based on pH change

• The change/shift in reaction based on the ions in the solution causes the indicator to change color from bonding/dissociating

• At the point of the color change, pKa = pH

Titration

• Neutralization reactions usually occur as titrations

• The pH increases slowly but then sharply after the equivalence point is reached, or when just enough base was added to neutralize the acid.

• At the center of the buffer region is the half-equivalence point, where just enough base was added to turn half of the acid to its conjugate base