Calculate the pH of a 1.35 M Solution of CH3CH2NH3Cl
Use this interactive calculator to determine the pH of ethylammonium chloride solution from the conjugate acid equilibrium of ethylamine. Adjust concentration, base constant, and calculation mode to explore the chemistry in detail.
Calculator
Ka = Kw / Kb
Expert Guide: How to Calculate the pH of a 1.35 M Solution of CH3CH2NH3Cl
To calculate the pH of a 1.35 M solution of CH3CH2NH3Cl, you need to recognize the chemical identity of the solute first. CH3CH2NH3Cl is ethylammonium chloride, the salt formed when the weak base ethylamine, CH3CH2NH2, reacts with hydrochloric acid. In water, the chloride ion acts as a spectator ion because it is the conjugate base of a strong acid and has negligible basicity under ordinary conditions. The important species is the ethylammonium ion, CH3CH2NH3+, which behaves as a weak acid. That means the pH of the solution is controlled by the acid dissociation equilibrium of CH3CH2NH3+ in water.
This type of problem appears frequently in general chemistry, analytical chemistry, and biochemistry because salts of weak bases are common laboratory reagents and industrial intermediates. The key to solving it correctly is to avoid treating the salt as neutral. Many students see a chloride salt and assume neutrality, but that only works for salts made from a strong acid and a strong base. Since ethylamine is a weak base, its conjugate acid is not neutral. It donates protons to water to a small but measurable extent, producing hydronium and lowering the pH below 7.
Step 1: Identify the acid species in solution
When ethylammonium chloride dissolves, it dissociates essentially completely:
CH3CH2NH3Cl → CH3CH2NH3+ + Cl-
The chloride ion does not significantly affect pH, so the relevant equilibrium is:
CH3CH2NH3+ + H2O ⇌ CH3CH2NH2 + H3O+
This is the weak acid equilibrium for the conjugate acid of ethylamine. Therefore, the pH calculation is a weak acid problem, not a strong acid problem.
Step 2: Convert Kb for ethylamine into Ka for ethylammonium
Since tables often list the base dissociation constant, Kb, for ethylamine rather than the acid dissociation constant, Ka, for ethylammonium, we use the standard conjugate relationship:
Ka × Kb = Kw
At 25 degrees C, Kw = 1.0 × 10-14. A common literature value for ethylamine is Kb = 5.6 × 10-4. Substituting gives:
Ka = (1.0 × 10^-14) / (5.6 × 10^-4) = 1.79 × 10^-11
Once Ka is known, the pKa is easy to find:
pKa = -log(Ka) ≈ 10.75
That high pKa confirms that CH3CH2NH3+ is a weak acid, but because the concentration is quite high at 1.35 M, enough hydronium still forms to make the solution appreciably acidic.
Step 3: Set up the ICE table
Let the initial concentration of CH3CH2NH3+ be 1.35 M, and let x be the amount that ionizes:
| Species | Initial (M) | Change (M) | Equilibrium (M) |
|---|---|---|---|
| CH3CH2NH3+ | 1.35 | -x | 1.35 – x |
| CH3CH2NH2 | 0 | +x | x |
| H3O+ | 0 | +x | x |
Substituting these into the equilibrium expression gives:
Ka = x^2 / (1.35 – x)
Step 4: Solve for x and then pH
Because Ka is small relative to the concentration, the weak acid approximation usually works well. If x is much smaller than 1.35, then 1.35 – x ≈ 1.35:
x ≈ √(KaC) = √[(1.79 × 10^-11)(1.35)] = 4.92 × 10^-6 M
Since x represents hydronium concentration:
[H3O+] ≈ 4.92 × 10^-6 M
Therefore:
pH = -log(4.92 × 10^-6) ≈ 5.31
If you solve the quadratic equation exactly, you get nearly the same value because x is tiny compared with 1.35 M. The approximation is valid here. A quick check shows:
(x / 1.35) × 100 ≈ 0.00036%
Since the percent ionization is far below 5%, the approximation is excellent.
Final answer for the default problem
For a 1.35 M solution of CH3CH2NH3Cl, using Kb = 5.6 × 10-4 for ethylamine at 25 degrees C, the calculated pH is approximately 5.31.
Why the pH is not extremely low despite a high concentration
A concentration of 1.35 M is large, so some learners expect a very acidic solution. The reason that does not happen is that CH3CH2NH3+ is a weak acid, not a strong acid. Only a very small fraction of the ions donate protons to water. In strong acid calculations, concentration and hydronium concentration are often nearly the same. In weak acid calculations, concentration provides the pool of acid, but the equilibrium constant controls how much of that pool actually ionizes. Here, the acid strength is so modest that only a few parts per million of the original concentration produce hydronium.
Comparison with related ammonium type ions
Comparing conjugate acids of common weak bases helps place this result into context. Smaller pKa means stronger conjugate acid and a lower pH at the same concentration. The table below shows typical values used in introductory chemistry references at 25 degrees C.
| Conjugate acid | Parent base | Typical Kb of base | Calculated Ka of conjugate acid | Typical pKa |
|---|---|---|---|---|
| NH4+ | NH3 | 1.8 × 10^-5 | 5.6 × 10^-10 | 9.25 |
| CH3NH3+ | Methylamine | 4.4 × 10^-4 | 2.3 × 10^-11 | 10.64 |
| CH3CH2NH3+ | Ethylamine | 5.6 × 10^-4 | 1.8 × 10^-11 | 10.75 |
| (CH3)2NH2+ | Dimethylamine | 5.4 × 10^-4 | 1.9 × 10^-11 | 10.72 |
These values reveal a useful pattern. Ethylammonium is a weaker acid than ammonium because ethylamine is a stronger base than ammonia. As a result, a concentrated ethylammonium chloride solution is acidic, but not as acidic as an equally concentrated ammonium chloride solution.
Estimated pH behavior at several concentrations
The concentration of the salt changes the pH, but not in a simple one to one linear fashion. For weak acids, hydronium concentration scales approximately with the square root of concentration when the weak acid approximation holds. The following estimates use Ka ≈ 1.79 × 10-11.
| CH3CH2NH3Cl concentration (M) | Estimated [H3O+] (M) | Estimated pH | Percent ionization |
|---|---|---|---|
| 0.010 | 4.23 × 10^-7 | 6.37 | 0.0042% |
| 0.100 | 1.34 × 10^-6 | 5.87 | 0.0013% |
| 1.00 | 4.23 × 10^-6 | 5.37 | 0.00042% |
| 1.35 | 4.92 × 10^-6 | 5.31 | 0.00036% |
Notice that raising concentration by a factor of 100 does not lower the pH by 2 full units as it would for a strong acid. That is one of the clearest signatures of weak acid behavior.
Common mistakes to avoid
- Assuming CH3CH2NH3Cl is neutral because it is a salt.
- Using HCl chemistry for the chloride ion. Chloride is a spectator here.
- Using Kb directly in the acid equilibrium instead of first converting it to Ka.
- Forgetting that CH3CH2NH3+ is the conjugate acid of a weak base.
- Using concentration as hydronium concentration. That would only work for a strong acid.
- Not checking whether the weak acid approximation is valid.
When should you use the exact quadratic equation?
For this problem, the approximation is more than adequate because Ka is very small and the concentration is large. However, using the exact quadratic solution is a good habit in software tools and advanced coursework because it removes guesswork. The exact equation comes from:
Ka = x^2 / (C – x)
Rearranging gives:
x^2 + Ka x – Ka C = 0
The physically meaningful root is:
x = (-Ka + √(Ka^2 + 4KaC)) / 2
This calculator includes both the exact and approximate methods so you can compare them directly. In the default case, the answers match to ordinary reporting precision, which reinforces that the chemistry is behaving exactly as expected.
Practical significance in laboratory and industrial settings
Ethylammonium salts appear in synthesis, separations, and formulation chemistry. A mildly acidic pH can influence reaction selectivity, extraction efficiency, corrosion behavior, and the protonation state of nearby compounds. Knowing how to calculate pH from equilibrium constants is therefore not just an academic exercise. It helps chemists predict how salts will behave in aqueous systems, especially when they are used in concentrated stock solutions. In process chemistry, a shift from pH 6.4 to pH 5.3 can matter for catalyst stability, decomposition pathways, and compatibility with pH sensitive materials.
Authoritative references for acid base constants and water equilibrium
Bottom line
The pH of a 1.35 M solution of CH3CH2NH3Cl is found by treating the ethylammonium ion as a weak acid. Start from the known Kb of ethylamine, convert it to Ka with Kw, set up the weak acid equilibrium, and solve for hydronium concentration. Using a typical Kb value of 5.6 × 10-4, the solution pH is about 5.31 at 25 degrees C. That result is chemically reasonable, mathematically consistent, and supported by standard acid base equilibrium theory.