Weak Acid & Base ICE Table Calculations 1412 2026 L21
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Overview
The lesson explains how to identify strong and weak acids and bases, represent their reactions in water, and calculate equilibrium concentrations and pH with ICE tables. It develops the shortcut that an initial concentration divided by Ka or Kb greater than 100 usually makes the small-x approximation acceptable, then applies it to acetic acid, ammonia, boric acid, amines, and cyanide salts, alongside Ka/Kb conversions, salt acidity, and percent ionization.
Key takeaways
- For weak-acid and weak-base ICE tables, compare the initial concentration with Ka or Kb: a ratio greater than 100 generally justifies dropping the −x term, with approximately 5% error at the boundary.
- At 25°C, conjugate acid and base constants satisfy KaKb = 1.0 × 10−14, so pKa + pKb = 14; this converts a supplied pKa into the Kb needed for a base calculation.
- Salt pH depends on the reactive ions: NH4+ is acidic, CN− is basic, and Na+ or NO3− is effectively neutral in the examples.
- Highly charged metal ions such as Al3+ and Cr3+ acidify water by polarizing coordinated water molecules and enabling proton donation.
- For 1.0 M acetic acid with Ka = 1.8 × 10−5, the small-x method gives [H3O+] ≈ 0.0042 M, pH ≈ 2.4, and percent ionization ≈ 0.42%.
- Use pH indicators according to their precision: phenolphthalein helps identify a titration endpoint, while pH paper is less reliable for distinguishing nearby values than an electrode.
Chapters
- The standard strong acids include HCl, HBr, HI, HNO3, HClO3, and HClO4; HF is weak because its small fluorine atom holds the H-F bond strongly.
- For H2SO4, the first proton dissociates essentially completely, while the second proton from HSO4− is weakly acidic.
- Organic acids such as acetic acid and benzoic acid are generally weak and require Ka or pKa data to compare strengths.
- A weak acid can be represented generically as HA + H2O ⇌ H3O+ + A−, while a base B accepts a proton from water to form BH+ and OH−.
- Highly charged cations such as Al3+ and Cr3+ coordinate water molecules; polarized O-H bonds allow a coordinated water molecule to donate H+, making the solution acidic.
- Soluble hydroxides such as NaOH dissociate completely; Ca(OH)2 produces Ca2+ and two OH− ions per formula unit.
- Bases without hydroxide, including methylamine and acetate, accept protons from water and generate OH−.
- The molecular, complete ionic, and net ionic forms of sodium acetate hydrolysis differ by whether ions are shown; Na+ cancels as a spectator ion.
- A larger Ka means a stronger weak acid; the lesson compares acetic acid, nitrous acid, and hydrogen sulfate, with hydrogen sulfate having the largest listed Ka.
- A larger Kb indicates a stronger base; among the examples, ammonia is stronger than acetate and nitrite.
- For a conjugate acid-base pair, Ka × Kb = Kw, approximately 1.0 × 10−14 at 25°C; equivalently, pKa + pKb = 14.
- Strong acids and bases dissociate almost completely in water, while weak acids and bases remain mostly as reactants at equilibrium.
- pH paper is portable and needs no batteries, but its color scale can make nearby values such as pH 9 and 10 difficult to distinguish; an electrode pH pen is more precise.
- Phenolphthalein turns pink after a strong-base titration passes its endpoint, so titrations should stop at the first persistent pale pink.
- Red cabbage extract changes color with pH, appearing red in acidic lemon juice and yellow in household ammonia.
- Butterfly pea flower tea can be served blue and shifted by lemon juice, providing a visible demonstration of pH-dependent color.
- For 1.0 M acetic acid with Ka = 1.8 × 10−5, the ICE table gives Ka = x²/(1 − x), where x is the equilibrium H3O+ concentration.
- When the initial concentration divided by Ka exceeds 100, the approximation 1 − x ≈ 1 is generally acceptable and gives x ≈ √(CinitialKa).
- The calculation gives approximately 4.2 × 10−3 M H3O+, pH 2.4, and pOH 11.6; Kw then gives approximately 2.4 × 10−12 M OH−.
- At the Cinitial/Ka = 100 boundary, the approximation produces about 5% error; larger ratios give smaller error.
- For aqueous ammonia, NH3 + H2O ⇌ NH4+ + OH−, use Kb rather than Ka in the equilibrium expression.
- Given pKa = 9.255 for NH4+, calculate pKb = 14 − 9.255 = 4.745, so Kb is approximately 1.8 × 10−5.
- For a 1.0 M NH3 solution, Cinitial/Kb is greater than 100, allowing the small-x approximation and producing about 4.2 × 10−3 M OH−.
- The resulting pOH is about 2.4 and pH about 11.6; recognizing that ammonia is a base is the key setup decision.
- Cr3+ and Al3+ make solutions acidic because their high charge polarizes coordinated water; the lesson also identifies other highly charged transition-metal ions as acidic.
- NO3−, the conjugate base of strong acid HNO3, has negligible effect on pH, while NH4+ acts as a weak acid.
- Na+ is effectively a spectator ion in water, whereas CN−, the conjugate base of weak acid HCN, accepts a proton and makes a cyanide salt solution basic.
- Classify each dissolved salt by identifying its cation and anion, then checking whether either ion is a weak acid, weak base, or spectator.
- For 3.1 M boric acid with Ka = 10−10, the large initial-concentration-to-Ka ratio supports the small-x approximation and gives a pH of about 4.4.
- Amines such as dimethylamine accept a proton from water to form a positively charged conjugate acid and OH−; complex bases such as lidocaine can be represented as B and BH+.
- For lidocaine, the example uses Cinitial/Kb = 74, applies the approximation, and obtains pH about 10.
- Conjugate species are formed by adding or removing H+; use KaKb = 10−14 to convert between a conjugate acid's Ka and a base's Kb.
- KCN dissociates to K+ and CN−; potassium is a spectator, while CN− hydrolyzes as a weak base.
- Given pKa = 9.21 for HCN, calculate pKb = 14 − 9.21 = 4.79, then use a base ICE table to find pH approximately 11.5 for the example solution.
- For 1.0 M acetic acid, an equilibrium ion concentration of about 0.0042 M corresponds to 0.42% ionization.
- Percent ionization is the equilibrium amount ionized divided by the initial acid concentration, multiplied by 100%.
Transcript excerpt
0:00 Yeah. [clears throat] Yeah. So, I made these for you guys. Um, there's always a fight for that proton between the bases and water is acting like a base. And if you have a weak acid, it's producing a conjugate base. And the conjugate base for the weak acid is way stronger than water. And so, I was looking on my little meme site for something with the big arm and the little arm and I couldn't find it. So, I went and I found a picture like that and I made this meme template just for this class. So, there you go. And then in terms of the strengths of acids, mineral acids like nitric acid and so on, those are way way way stronger than any organic acid. And so that's what I'm trying to show with that one. So organic acids, anything we can with that CO that you see or just a bunch of carbons and
0:48 hydrogens and oxygens, an organic molecule as an acid, that's going to be a weak acid. you need to look up the the the Ka or the pKa for that acid to know how strong it is. So, we're going to be talking about weak acids and bases today. Um, [clears throat] we'll start out by saying these are not weak. These are all of your strong acids. So, perchloric and hydrochloric and hydroic and hydroiotic all of those down the hallogen row there. Now, notice what's not on there is a strong acid, the top one, hydrofuloric acid. So hydrofuloric acid is a weak acid and it's because that florine is such a small atom that it's really wrestling with that hydrogen and they'd rather stay together. It's more of a covealent bond. But when you get to the bigger chlorine, bromine,
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