Calculate OH-, H+ and the pH of 20 M Triethylamine
Use this premium weak-base calculator to estimate hydroxide concentration, hydrogen ion concentration, pOH, and pH for triethylamine solutions. The default setup is prefilled for 20.0 M triethylamine at 25 degrees Celsius, using a typical pKb value near 3.25.
Triethylamine pH Calculator
This tool solves the weak-base equilibrium exactly with the quadratic expression rather than relying only on the small-x approximation.
Press Calculate to see OH-, H+, pOH, pH, percent ionization, and a chart of the equilibrium result.
Expert guide: how to calculate OH-, H+, and the pH of 20 M triethylamine
Triethylamine is a classic weak organic base used in synthesis, separations, catalysis support, and pH adjustment in laboratory work. When students or researchers ask how to calculate OH-, H+, and the pH of 20 M triethylamine, they are really asking for a weak-base equilibrium calculation. Unlike a strong base such as sodium hydroxide, triethylamine does not dissociate completely in water. Instead, it reacts reversibly with water to produce triethylammonium and hydroxide:
(C2H5)3N + H2O ⇌ (C2H5)3NH+ + OH-
The key constant for this process is the base dissociation constant, Kb, or its logarithmic form, pKb. For triethylamine, a commonly cited pKb at room temperature is about 3.25, corresponding to Kb around 5.62 × 10-4. Once Kb and the starting concentration are known, you can determine the equilibrium hydroxide concentration. From there, pOH, pH, and hydrogen ion concentration follow directly. For a 20 M solution, the pH is strongly basic, but because triethylamine is weak, the hydroxide concentration is still far below 20 M.
Why triethylamine requires an equilibrium calculation
A common mistake is to treat triethylamine as if it were a strong base. If that were true, a 20 M solution would give an OH- concentration close to 20 M and a pH far outside realistic dilute-solution assumptions. In reality, weak bases establish equilibrium and only a fraction of the molecules accept a proton from water. That means the actual OH- concentration must be found from the equilibrium expression:
Kb = [BH+][OH-] / [B]
Let the starting concentration be C and let x be the amount that reacts. Then:
- [B] at equilibrium = C – x
- [BH+] at equilibrium = x
- [OH-] at equilibrium = x
Substituting these values gives:
Kb = x2 / (C – x)
For a quick estimate, many textbooks assume x is small compared with C, which produces x ≈ √(KbC). That approximation is usually acceptable when percent ionization is low. However, this calculator also offers the exact quadratic solution, which is the better choice whenever you want a cleaner answer.
Step by step calculation for 20 M triethylamine at 25 degrees C
- Start with concentration, C = 20.0 M.
- Use pKb = 3.25, so Kb = 10-3.25 ≈ 5.62 × 10-4.
- Set up the equilibrium expression: Kb = x2 / (20.0 – x).
- Solve exactly: x = (-Kb + √(Kb2 + 4KbC)) / 2.
- This x value is the equilibrium [OH-].
- Compute pOH = -log10[OH-].
- At 25 degrees C, use pH = 14.00 – pOH.
- Finally, [H+] = 10-14 / [OH-], or equivalently 10-pH.
Using these inputs, the equilibrium OH- concentration is about 0.1057 M. That gives a pOH of about 0.976 and a pH near 13.024 at 25 degrees C. The corresponding hydrogen ion concentration is about 9.46 × 10-14 M. Even though the formal triethylamine concentration is 20 M, only about 0.53% ionizes under this simple weak-base model. This is exactly why weak-base chemistry matters: concentration alone does not tell you pH unless you also know the acid-base strength.
| Parameter | Value for 20 M triethylamine | Notes |
|---|---|---|
| pKb | 3.25 | Typical textbook value near room temperature |
| Kb | 5.62 × 10-4 | Computed from 10-pKb |
| [OH-] | 0.1057 M | Exact quadratic equilibrium solution |
| pOH | 0.976 | -log10[OH-] |
| pH | 13.024 | At 25 degrees C, pH = 14.00 – pOH |
| [H+] | 9.46 × 10-14 M | From Kw / [OH-] |
| Percent ionization | 0.53% | 100 × [OH-] / initial concentration |
Approximation versus exact solution
Many chemistry problems are designed so the approximation x ≈ √(KbC) works. For 20 M triethylamine, that estimate gives x ≈ √(5.62 × 10-4 × 20) ≈ 0.1060 M, which is very close to the exact value. The difference is tiny because the ionization remains low relative to the formal concentration. Still, exact methods are preferable in a calculator because they avoid avoidable error and remain reliable across a broader range of inputs.
| Method | Calculated [OH-] | pH at 25 degrees C | Relative difference |
|---|---|---|---|
| Exact quadratic | 0.1057 M | 13.024 | Reference |
| Approximation, √(KbC) | 0.1060 M | 13.025 | Less than 0.3% in [OH-] |
| Incorrect strong-base assumption | 20.0 M | Not chemically appropriate in this model | Massively overestimates OH- |
Important chemistry context for concentrated solutions
There is an important practical caveat. A nominal 20 M triethylamine solution is extremely concentrated. In highly concentrated real solutions, ideal-solution assumptions break down. Activity effects, density changes, nonideal interactions, and solvent limitations can make the true experimental pH depart from the simple equilibrium value calculated from molarity alone. In other words, the weak-base equation is the standard educational method and a useful estimate, but not the last word in precision physical chemistry for very concentrated systems.
That is why professional chemists distinguish between concentration and activity. Most introductory and intermediate pH problems use concentration-based equilibrium expressions. Those are excellent for learning and for many routine calculations. But if you are doing advanced formulation, electrochemistry, or calibration-sensitive analytical work, you should check whether activity coefficients, ionic strength corrections, or direct pH measurement are required.
How OH-, H+, pOH, and pH are connected
- [OH-] tells you the hydroxide ion concentration produced by base hydrolysis.
- pOH is the negative base-10 logarithm of [OH-].
- pH is related to pOH by pH + pOH = pKw. At 25 degrees C, pKw is 14.00.
- [H+] comes from water autoionization: [H+][OH-] = Kw.
Once one of these values is known, the others can be found quickly. In a triethylamine problem, [OH-] is usually the first quantity you solve for because the compound is a base. Then the rest follow from logarithms and the water ion product.
Common mistakes to avoid
- Using Ka instead of Kb for the base itself.
- Forgetting to convert pKb to Kb with 10-pKb.
- Treating triethylamine as a strong base.
- Using pH = 14 – pOH without checking the temperature assumption.
- Reporting too many digits when the input data are approximate.
- Ignoring that highly concentrated solutions can behave nonideally.
What the calculator on this page does
The calculator above lets you enter the formal concentration, pKb, temperature setting, and your preferred numerical method. On click, it reads the inputs, computes Kb, solves the equilibrium, and returns the estimated hydroxide concentration, hydrogen ion concentration, pOH, pH, and percent ionization. It also plots a chart so you can visually compare the starting base concentration with the amount converted to hydroxide and conjugate acid.
This can be especially useful when comparing different weak bases, studying the effect of concentration, or checking whether the approximation remains valid across a range of conditions. For triethylamine specifically, it helps demonstrate a key concept: a weak base can still produce a high pH if the solution is concentrated enough, even when only a small fraction ionizes.
Authoritative references for acid-base data and water chemistry
For additional background and reference material, review authoritative educational and government resources:
- LibreTexts Chemistry for broad acid-base equilibrium explanations.
- U.S. Environmental Protection Agency for water chemistry context and pH fundamentals.
- NIST Chemistry WebBook for reference chemical information and data context.
Bottom line
To calculate OH-, H+, and the pH of 20 M triethylamine, treat triethylamine as a weak base, not a strong one. Use its Kb or pKb, solve the base-hydrolysis equilibrium, then convert [OH-] to pOH and pH. With pKb around 3.25 at 25 degrees C, a 20 M solution gives an estimated [OH-] near 0.1057 M, a pOH near 0.976, a pH near 13.024, and an [H+] around 9.46 × 10-14 M. Those values are the standard concentration-based answer for chemistry calculations, while remembering that very concentrated real solutions may show nonideal behavior in experimental settings.