CHEM1312H - ICF vs. ICE and Types of Titration Problems

ICF chart(initial, change, and final)

a chart or stoichiometry table used to determine a limiting reactant and can only be used when the reaction is NOT at equilibrium (moles should only be included in ICF chart) (ex: making a buffer by reacting a strong acid with a weak base (or basic salt) or a strong base and a weak acid (or acidic salt)

when to use an ICF chart

-addition of a small amount of strong acid or strong base to a buffer

-titration of a strong acid with a strong base or a strong base with a strong acid

-titration of a weak acid with a strong base or a weak base with a strong acid

ICE (initial, change, and equilibrium)

-chart used to determine the OH^- or H₃O^- ion concentration for a substance that does not completely ionize

-includes only concentrations in M (mol/L)

when to use an ICE chart

-salts that yield acidic or basic solutions

-these salts are strong electrolytes, and they will completely dissociate in water and the concentration of the ion undergoing hydrolysis in water will correspond to the concentration of the salt

types of acid-base reactions

-reaction of a strong acid and strong base

-reaction of a weak acid and strong base

-reaction of a strong acid and weak base

-reaction of a weak acid and weak base

reaction of a strong acid and strong base

-net ionic equation for strong acid reacting with strong base is always: H⁺(aq) + OH^-(aq) → H₂O(l)

-pH of resultant reaction is always 7

reaction of a weak acid and strong base

hydrolysis of anion, pH > 7

reaction of a strong acid and weak base

hydrolysis of cation, pH < 7

reaction of a weak acid and weak base

-weak acid and base neutralization: pH could be 7, >7, or <7, depending on the K_a of weak acid and K_b of weak base

-neutralization products almost neutral: pH ~ 7

if you are titrating a strong acid with strong base (or strong base with strong acid)

-the pH before addition of strong base is the pH of the strong acid solution. use pH = -log[H₃O⁺] or pOH = -log[OH^-] in case of a strong base titration and then use pH + pOH = 14 to determine the pH

-after addition of strong base, use an ICF table to determine the limiting reactant. if strong acid is limiting reactant, then the strong base is remaining and the strong base is limiting reactant, then strong acid is remaining

-pH at equivalence point is 7 (moles of H₃O⁺ = moles of OH^-)

if you are titrating a weak acid with a strong base

-pH before addition of strong base is the pH of the weak acid, use an ICE table to determine [H₃O⁺]

-pH after addition of strong base, use an ICF table to determine the limiting reactant

-if both the weak acid and strong base are limiting reactants (equivalence point), then you will be left with a basic salt, ICE chart to determine [OH^-] and then pH (pH > 7)

if you are titrating a weak acid with strong base: limiting reactants

-if strong base is limiting reactant, you will be left with weak acid and salt of conjugate base, buffer and use henderson-hasselbalch equation to calculate the pH (pH = pK_a + log[A^-]/[HA]

-if weak acid is limiting reactant, then strong base will remain (ignore the weak base if it exists with a strong base due to common-ion effect)

if you are titrating a weak base with strong acid

-pH before addition of strong acid is the pH of the weak base, use an ICE table to determine [OH^-] and then determine pOH and use pOH + pH = 14 to determine the pH

-if both the weak base and strong acid are limiting reactants (equivalence point), then you will be left with an acidic salt, ICE chart to determine (H₃O⁺] and then pH (pH < 7)

if you are titrating a weak base with a strong acid: limiting reactant

-if strong acid the limiting reactant, then you will be left with weak base and salt of conjugate acid, buffer and use HH equation to calculate pOH to determine pH

-if weak base is limiting reactant, then strong acid will only remain (ignore the weak acid salt if it exists withs a strong acid due to common-ion effect