Calculate pH of 0.10 M C4H9NH2 Measured at 12.04
Use this interactive chemistry calculator to analyze a 0.10 M butylamine solution, confirm the measured pH of 12.04, and derive pOH, hydroxide concentration, hydronium concentration, percent ionization, Kb, and pKb from the observed data.
This calculator treats the measured pH as the experimental value and back-calculates equilibrium quantities at 25 degrees C using pH + pOH = 14.
Results will appear here
Click Calculate to analyze the 0.10 M C4H9NH2 solution measured at pH 12.04.
How to Calculate the pH of 0.10 M C4H9NH2 Measured at 12.04
When a chemistry problem asks you to calculate the pH of 0.10 M C4H9NH2 measured at 12.04, the key idea is that the solution has already been experimentally observed. In other words, the pH is not merely theoretical here. The pH meter reading of 12.04 is the central piece of data, and from it you can derive several important equilibrium quantities for butylamine, a weak base. C4H9NH2, commonly called butylamine, reacts with water to produce hydroxide ions and its conjugate acid. Because it is a weak base rather than a strong base, it does not ionize completely, which is why an equilibrium treatment is the correct analytical approach.
For students, lab users, and anyone checking an answer key, the fastest conclusion is simple: the pH of the solution is 12.04. However, most chemistry assignments want more than that. They typically want you to interpret the meaning of that measured pH. That means calculating pOH, determining the hydroxide concentration, estimating the concentration of protonated butylamine, calculating the percent ionization of the base, and often deriving an experimental value of the base dissociation constant Kb.
Step 1: Write the weak base equilibrium reaction
Butylamine behaves as a Brønsted base in water:
C4H9NH2 + H2O ⇌ C4H9NH3+ + OH-
This reaction tells you that each mole of butylamine that accepts a proton creates one mole of hydroxide ion. That one-to-one stoichiometric relationship is what makes the pH measurement so useful. Once you know the hydroxide concentration, you can infer how much butylamine has reacted.
Step 2: Convert pH to pOH
At 25 degrees C, the relationship between pH and pOH is:
pH + pOH = 14.00
Given pH = 12.04:
pOH = 14.00 – 12.04 = 1.96
This immediately confirms the solution is basic, because the pOH is low and the pH is well above 7.
Step 3: Find the hydroxide ion concentration
Use the definition of pOH:
[OH-] = 10^(-pOH)
Substitute the value:
[OH-] = 10^(-1.96) = 1.10 x 10^-2 M
This value is the equilibrium hydroxide concentration in the solution. Because the reaction produces hydroxide and protonated butylamine in equal amounts, this also means:
- [C4H9NH3+] at equilibrium is approximately 1.10 x 10^-2 M
- The amount of unreacted C4H9NH2 is the initial concentration minus that ionized amount
Step 4: Calculate the hydronium concentration
Use the definition of pH:
[H3O+] = 10^(-pH)
Substituting the measured value:
[H3O+] = 10^(-12.04) = 9.12 x 10^-13 M
This very small hydronium concentration is exactly what you expect from a strongly basic aqueous solution.
Step 5: Determine equilibrium concentrations
Start with an ICE setup:
- Initial [C4H9NH2] = 0.10 M
- Initial [C4H9NH3+] = 0
- Initial [OH-] is often taken as 0 for the weak-base equilibrium setup because water contributes only a negligible amount compared with the measured basicity
If x is the amount ionized, then at equilibrium:
- [C4H9NH2] = 0.10 – x
- [C4H9NH3+] = x
- [OH-] = x
From the measured pH, we already found x = 1.10 x 10^-2 M. Therefore:
- [C4H9NH2]eq = 0.10 – 0.0110 = 0.0890 M
- [C4H9NH3+]eq = 0.0110 M
- [OH-]eq = 0.0110 M
| Quantity | Value for 0.10 M C4H9NH2 at pH 12.04 | Why it matters |
|---|---|---|
| pH | 12.04 | Direct experimental acidity-basicity measurement |
| pOH | 1.96 | Lets you calculate hydroxide concentration |
| [OH-] | 1.10 x 10^-2 M | Represents base-generated hydroxide at equilibrium |
| [H3O+] | 9.12 x 10^-13 M | Shows hydronium is extremely low in this basic solution |
| [C4H9NH3+] | 1.10 x 10^-2 M | Conjugate acid formed from butylamine |
| [C4H9NH2] remaining | 8.90 x 10^-2 M | Unreacted weak base after equilibrium is reached |
| Percent ionization | 11.0% | Indicates a meaningful but incomplete ionization typical of a weak base |
Step 6: Calculate percent ionization
Percent ionization tells you how much of the weak base reacted:
Percent ionization = (x / initial concentration) x 100
= (0.0110 / 0.10) x 100 = 11.0%
This is an important interpretation point. A strong base at the same nominal concentration would ionize nearly 100%, but butylamine only ionizes partially. That is why weak bases require equilibrium analysis instead of simple complete-dissociation assumptions.
Step 7: Calculate the experimental Kb
The base dissociation expression is:
Kb = [C4H9NH3+][OH-] / [C4H9NH2]
Plug in the equilibrium concentrations:
Kb = (0.0110 x 0.0110) / 0.0890
Kb ≈ 1.36 x 10^-3
Then:
pKb = -log(Kb) ≈ 2.87
This experimentally derived Kb is somewhat stronger than many tabulated values often reported for normal butylamine in introductory chemistry references. That difference can happen for several reasons, including temperature drift, electrode calibration, ionic strength, concentration rounding, activity effects, contamination, or differences among butylamine isomers.
Why the measured pH can differ from a textbook estimate
A common classroom method is to start with a literature Kb value and predict the pH of a 0.10 M weak base solution. For butylamine, many tables place Kb in the neighborhood of 4 x 10^-4, which would produce a theoretical pH somewhat below 12.04. Yet your measured value is 12.04, and in experimental chemistry the measured data should guide the analysis if the problem explicitly gives it. This is a useful lesson: theory predicts, but measurement decides.
Several practical effects can shift weak-base pH measurements:
- Temperature: pH and pOH relationships are temperature dependent, although many classroom problems assume 25 degrees C.
- Probe calibration: A small offset in a pH meter has a noticeable effect on logarithmic calculations.
- Activity versus concentration: Real solutions do not always behave ideally, especially as ionic strength increases.
- Purity and isomer identity: C4H9NH2 can refer broadly to butylamine formulas, and different isomers can show somewhat different behavior.
- Rounding: Entering 0.10 M and 12.04 already limits precision.
Comparison data: how butylamine stacks up against other common weak bases
The following comparison table gives useful real-world chemistry context. The values below are typical approximate literature-level values at room temperature used in general chemistry for comparison. Exact numbers vary slightly by source and conditions, but the trends are reliable: lower pKb means a stronger base and generally a higher pH at the same formal concentration.
| Weak base | Approximate Kb | Approximate pKb | Estimated pH of 0.10 M solution |
|---|---|---|---|
| Ammonia, NH3 | 1.8 x 10^-5 | 4.75 | 11.13 |
| Methylamine, CH3NH2 | 4.4 x 10^-4 | 3.36 | 11.82 |
| Ethylamine, C2H5NH2 | 6.4 x 10^-4 | 3.19 | 11.90 |
| Butylamine, C4H9NH2 | about 4.0 x 10^-4 | about 3.40 | about 11.80 |
| Experimental butylamine from this measured pH | 1.36 x 10^-3 | 2.87 | 12.04 observed |
Common mistakes when solving this problem
- Using pH directly as [H3O+]: pH is a logarithm, not a concentration.
- Forgetting to convert pH to pOH: Weak bases are easiest to analyze through hydroxide concentration.
- Assuming complete dissociation: Butylamine is a weak base, so it does not fully ionize.
- Ignoring the wording “measured at 12.04”: That phrase means experimental pH is given and should be used, not replaced by a purely theoretical estimate.
- Dropping units: Concentration values should be reported in molarity.
- Over-rounding intermediate calculations: Keep extra digits until the final answer.
What this result says chemically
A pH of 12.04 for a 0.10 M butylamine solution tells you several things about the chemistry of the sample. First, the solution is distinctly basic, far above neutral water. Second, enough hydroxide is generated to make the solution noticeably alkaline, but not enough to suggest complete ionization. Third, butylamine has a meaningful affinity for protons, yet it remains in equilibrium with its conjugate acid rather than behaving like sodium hydroxide or potassium hydroxide. This is the hallmark of an organic amine base in water.
In practical laboratory terms, this means the solution can neutralize acids, alter indicator color, and affect reaction pathways in synthesis or titration work. If you are analyzing a weak-base sample, the pH measurement also provides a fast way to estimate whether the material concentration, purity, or identity is consistent with expectations. That is why pH-based back-calculation is such an important technique in analytical chemistry and chemical education.
Authority sources for pH and butylamine data
For deeper study, these authoritative sources are useful references:
- U.S. Environmental Protection Agency: pH overview and interpretation
- NIST Chemistry WebBook: butylamine compound data
- University of Wisconsin chemistry tutorial on weak bases and equilibrium concepts
Final takeaway
If you need to calculate the pH of 0.10 M C4H9NH2 measured at 12.04, the direct answer is that the pH is 12.04 because it is experimentally given. The more complete chemistry answer is that this corresponds to pOH = 1.96, hydroxide concentration of 1.10 x 10^-2 M, hydronium concentration of 9.12 x 10^-13 M, equilibrium butylammonium concentration of 1.10 x 10^-2 M, remaining butylamine concentration of 8.90 x 10^-2 M, percent ionization of 11.0%, and an experimental Kb of approximately 1.36 x 10^-3. If your instructor or lab asks for interpretation, explain that butylamine is a weak base and that the measured pH gives a direct route to the equilibrium composition of the solution.