Calculate The Ph Of 0.35 M Ethylamine

Use this premium weak-base calculator to find pOH, pH, hydroxide concentration, and equilibrium composition for aqueous ethylamine at 25 degrees Celsius.

Ethylamine pH Calculator

Expert Guide: How to Calculate the pH of 0.35 M Ethylamine

Ethylamine is a classic weak base problem in general chemistry. When you are asked to calculate the pH of 0.35 M ethylamine, the central idea is that ethylamine does not fully ionize in water. Instead, it partially reacts with water to produce ethylammonium ions and hydroxide ions. Because pH depends on hydrogen ion concentration and weak bases are often easier to solve through hydroxide production, the standard path is to calculate the hydroxide concentration first, determine pOH, and then convert that to pH.

Ethylamine, often written as C₂H₅NH₂, behaves as a Brønsted base because the nitrogen atom has a lone pair of electrons that can accept a proton from water. In aqueous solution, the relevant equilibrium is:

C₂H₅NH₂ + H₂O ⇌ C₂H₅NH₃⁺ + OH^–

That equilibrium is governed by the base dissociation constant, Kb. For ethylamine at about 25 degrees Celsius, a commonly used textbook value is around 5.6 × 10^-4. Since that Kb is much smaller than 1, the reaction is product-favored only to a limited extent, which is exactly why ethylamine is treated as a weak base rather than a strong base.

Step 1: Identify the Known Values

• Initial ethylamine concentration, C = 0.35 M
• Base dissociation constant, Kb = 5.6 × 10^-4
• Need to find pH

At the start, before equilibrium, we assume the hydroxide ion concentration generated by ethylamine is negligible compared with the amount that will form from the weak-base reaction. We also typically neglect the tiny 1.0 × 10^-7 M hydroxide that comes from water autoionization because it is insignificant relative to the amount produced by a 0.35 M weak base.

Step 2: Build the ICE Table

An ICE table is the most organized way to approach this problem.

Species	Initial (M)	Change (M)	Equilibrium (M)
C2H5NH2	0.35	-x	0.35 – x
C2H5NH3^+	0	+x	x
OH^–	0	+x	x

The equilibrium expression is:

Kb = [C₂H₅NH₃⁺][OH^–] / [C₂H₅NH₂]

Substituting the ICE table values:

5.6 × 10^-4 = x² / (0.35 – x)

Step 3: Solve for x

There are two acceptable ways to solve this type of weak-base problem: the approximation method and the quadratic method. In educational settings, the approximation method is often taught first, but the quadratic method is the more exact route and is what this calculator uses by default.

Approximation Method

If x is small relative to 0.35, then 0.35 – x is approximately 0.35. The equilibrium equation becomes:

5.6 × 10^-4 ≈ x² / 0.35

So:

x² ≈ (5.6 × 10^-4)(0.35) = 1.96 × 10^-4

Taking the square root:

x ≈ 1.40 × 10^-2 M

Since x represents [OH^–] at equilibrium, the hydroxide concentration is about 0.0140 M.

Now calculate pOH:

pOH = -log(0.0140) ≈ 1.85

Then convert to pH:

pH = 14.00 – 1.85 = 12.15

Quadratic Method

For higher confidence, especially if your instructor requires exact work, solve:

5.6 × 10^-4 = x² / (0.35 – x)

Rearranging gives:

x² + (5.6 × 10^-4)x – (1.96 × 10^-4) = 0

Using the quadratic formula yields a positive root very close to 0.01372 M. This leads to:

• [OH^–] ≈ 0.01372 M
• pOH ≈ 1.86
• pH ≈ 12.14

So the pH of 0.35 M ethylamine is approximately 12.14 to 12.15, depending on the chosen Kb value and whether rounding is performed during the intermediate steps.

Why the Answer Is Basic

Because ethylamine is a base, it raises the hydroxide concentration above that of pure water. Any solution with pH above 7 is basic at 25 degrees Celsius. Here the pH is a little above 12, which means the solution is distinctly basic, though still not as extreme as a strong base of similar formal concentration. That difference matters. A 0.35 M strong base would generate much more hydroxide than a weak base like ethylamine because strong bases dissociate essentially completely.

Solution	Concentration (M)	Model Used	Approximate [OH^–] (M)	Approximate pH at 25 degrees Celsius
Ethylamine	0.35	Weak base, Kb = 5.6 × 10^-4	0.0137	12.14
NaOH	0.35	Strong base, full dissociation	0.35	13.54
NH3	0.35	Weak base, Kb ≈ 1.8 × 10^-5	0.00251	11.40

This comparison table shows that ethylamine is a stronger weak base than ammonia, but still far weaker than sodium hydroxide. The result is chemically sensible: the pH for 0.35 M ethylamine should fall between a strong base at the same concentration and a weaker base such as ammonia.

Percent Ionization of Ethylamine

Another useful quantity is percent ionization, which tells you what fraction of the original base molecules actually reacted with water.

Percent ionization = (x / 0.35) × 100

Using x ≈ 0.01372 M:

Percent ionization ≈ (0.01372 / 0.35) × 100 ≈ 3.92%

That number explains why the approximation method works fairly well here. Since less than 5% of the original concentration is consumed, the change is modest enough that replacing 0.35 – x with 0.35 introduces only a small error.

What Students Often Get Wrong

1. Using Ka instead of Kb. Ethylamine is a base, so Kb is the correct equilibrium constant unless you are explicitly given the conjugate acid and asked to work backward.
2. Forgetting the pOH step. A weak base problem usually gives you hydroxide, not hydronium, so calculate pOH first and then convert to pH.
3. Treating ethylamine as a strong base. Do not assume [OH^–] = 0.35 M. That would significantly overestimate the pH.
4. Dropping x without checking. The approximation should be checked with percent ionization or by confirming that x is small relative to the initial concentration.
5. Rounding too early. Keep extra digits until the final pH value to avoid small but noticeable shifts in the answer.

Why Kb Values Can Vary Slightly by Source

You may notice that different textbooks or databases list slightly different Kb values for ethylamine. That is normal. Reported equilibrium constants can differ because of temperature, ionic strength, rounding, and the reference source used. A small shift in Kb slightly changes the final pH, but not enough to change the conclusion that a 0.35 M ethylamine solution is strongly basic. For example, a Kb near 4.3 × 10^-4 would produce a pH a little lower, while a Kb near 6.4 × 10^-4 would produce a pH a little higher.

Kb Assumed	Calculated [OH^–] (M)	Calculated pH	Interpretation
4.3 × 10^-4	0.01206	12.08	Basic, slightly lower estimate
5.6 × 10^-4	0.01372	12.14	Common textbook estimate
6.4 × 10^-4	0.01464	12.17	Basic, slightly higher estimate

Fast Mental Check for Reasonableness

If you want a quick confidence check, think in ranges. A strong base at 0.35 M would have pH around 13.54. Ammonia at the same concentration gives a lower pH, around 11.4. Ethylamine is a stronger base than ammonia but still weak compared with sodium hydroxide, so a pH in the low 12s is perfectly plausible. That kind of estimation is valuable during exams because it helps you catch calculator mistakes immediately.

Authoritative Chemistry References

If you want to validate weak-base concepts, pH theory, and equilibrium calculations from authoritative educational or government-backed sources, these references are useful:

• LibreTexts Chemistry educational resource
• U.S. Environmental Protection Agency overview of pH
• NIST Chemistry WebBook

Final Answer

Using a typical value of Kb = 5.6 × 10^-4 for ethylamine at 25 degrees Celsius, the calculated hydroxide concentration for a 0.35 M solution is approximately 0.0137 M. That gives a pOH of about 1.86 and a final pH of about 12.14.

Bottom line: The pH of 0.35 M ethylamine is approximately 12.14 under standard textbook conditions at 25 degrees Celsius.