Calculate The Ph Of A 0100M Ethylamine Solution If

This premium calculator helps you determine the pH, pOH, hydroxide concentration, conjugate acid concentration, and percent ionization for an aqueous ethylamine solution. Enter the concentration, choose whether you want to use Kb or pKb, and the calculator will solve the weak-base equilibrium exactly.

Ethylamine pH Calculator

Default values reflect a common textbook case: 0.100 M ethylamine at 25 degrees Celsius, using a typical base dissociation constant near 5.6 × 10^-4.

Model used:

For the weak base equilibrium C2H5NH2 + H2O ⇌ C2H5NH3+ + OH-, the calculator solves Kb = x² / (C – x) exactly with the quadratic expression x = (-Kb + √(Kb² + 4KbC)) / 2, where x = [OH-].

Expert Guide: How to Calculate the pH of a 0.100 M Ethylamine Solution

If you are asked to calculate the pH of a 0.100 M ethylamine solution, you are working with a classic weak-base equilibrium problem from general chemistry. Ethylamine, written as C₂H₅NH₂, is an amine that accepts a proton from water rather than donating one. That means it raises the hydroxide concentration, produces a basic solution, and requires a base equilibrium approach rather than a strong-base shortcut.

The wording of these problems often ends with “if” because the full question usually provides one more key fact such as the Kb of ethylamine, the pKb, or a temperature assumption. Once you know the concentration and the equilibrium constant, the path to pH is straightforward: write the chemical equation, set up an ICE table, solve for the hydroxide concentration, calculate pOH, and finally convert pOH into pH.

Ethylamine is a weak base, which means it does not react completely with water. Instead, it partially ionizes according to the equilibrium:

C2H5NH2(aq) + H2O(l) ⇌ C2H5NH3+(aq) + OH-(aq)

Because the base is weak, the amount that reacts is only a fraction of the original 0.100 M concentration. This is why we cannot simply say the hydroxide concentration is 0.100 M. We must calculate the equilibrium amount. For many textbook sets, a commonly cited Kb for ethylamine at 25 degrees Celsius is around 5.6 × 10^-4, though some tables report values very close to that depending on rounding and source.

Step-by-Step Method

1. Write the balanced base ionization equation for ethylamine in water.
2. Set up an ICE table using the initial concentration of 0.100 M.
3. Express Kb in terms of the unknown change, usually called x.
4. Solve for x, where x equals the equilibrium hydroxide concentration.
5. Find pOH using pOH = -log[OH-].
6. Find pH using pH = 14.00 – pOH at 25 degrees Celsius.

ICE Table for 0.100 M Ethylamine

Let the initial ethylamine concentration be 0.100 M. Assume no significant initial concentration of ethylammonium ion or hydroxide from the base itself.

Species	Initial (M)	Change (M)	Equilibrium (M)
C2H5NH2	0.100	-x	0.100 – x
C2H5NH3^+	0	+x	x
OH^–	0	+x	x

Plug those equilibrium expressions into the base dissociation constant:

Kb = [C2H5NH3+][OH-] / [C2H5NH2] = x² / (0.100 – x)

If we use Kb = 5.6 × 10^-4, then:

5.6 × 10^-4 = x² / (0.100 – x)

Solving this exactly gives x, the equilibrium hydroxide concentration, at about 0.00722 M. Then:

• [OH-] ≈ 0.00722 M
• pOH = -log(0.00722) ≈ 2.141
• pH = 14.000 – 2.141 = 11.859

So the pH of a 0.100 M ethylamine solution is approximately 11.86 when Kb is taken as 5.6 × 10^-4 at 25 degrees Celsius.

Using pKb Instead of Kb

Some instructors or problem sets give pKb instead of Kb. The relationship is:

Kb = 10^(-pKb)

If pKb = 3.25, then:

Kb = 10^(-3.25) ≈ 5.62 × 10^-4

That leads to nearly the same pH result. This is why your final answer may differ slightly depending on whether the source rounded Kb or pKb first.

Approximation Versus Exact Solution

In many weak acid and weak base problems, students are taught the approximation C – x ≈ C. For a 0.100 M ethylamine solution, the approximation works fairly well because x is much smaller than 0.100, although it is not vanishingly small. If we use the approximation:

x ≈ √(Kb × C) = √(5.6 × 10^-4 × 0.100) ≈ 0.00748 M

That gives pOH ≈ 2.126 and pH ≈ 11.874. The difference from the exact quadratic answer is small, about 0.015 pH units. In classroom work, this may be acceptable depending on the required precision, but in an online calculator or a rigorous lab report, the exact method is better.

Method	[OH^–] (M)	pOH	pH	Difference from Exact
Exact quadratic solution	0.00722	2.141	11.859	Reference value
Square-root approximation	0.00748	2.126	11.874	+0.015 pH units

Why Ethylamine Is Basic

Ethylamine contains a nitrogen atom with a lone pair. That lone pair can accept a proton from water, producing the conjugate acid ethylammonium and hydroxide ions. Organic amines are often weak bases because proton transfer is favorable but not complete. Compared with ammonia, many alkyl amines are somewhat stronger bases in water because alkyl groups can donate electron density toward nitrogen, increasing proton affinity.

This idea also helps explain why the pH of a 0.100 M ethylamine solution is clearly basic but still far below the pH expected for a 0.100 M strong base such as sodium hydroxide.

Comparison with Other Common Bases

The table below compares representative basicity data at 25 degrees Celsius. Exact values can vary slightly among references, but the listed statistics are consistent with widely used chemistry tables and instructional materials.

Base	Representative Kb at 25 degrees C	Representative pKb	Expected Behavior in Water
Ethylamine	5.6 × 10^-4	3.25	Weak base, partially ionizes
Ammonia	1.8 × 10^-5	4.74	Weaker base than ethylamine
Aniline	4.3 × 10^-10	9.37	Very weak base in water
Sodium hydroxide	Complete dissociation	Not treated with Kb in intro courses	Strong base, essentially full ionization

Common Mistakes to Avoid

• Using pH = -log(0.100). That would only work for a strong acid, not a weak base.
• Forgetting that ethylamine generates OH^–, so you find pOH first and then convert to pH.
• Using Ka instead of Kb. Ethylamine is a base, so Kb is the direct equilibrium constant.
• Ignoring temperature assumptions. The common formula pH + pOH = 14 applies at 25 degrees Celsius when pKw = 14.00.
• Applying the approximation without checking whether it is reasonable.

When the Problem Says “If”

Textbook prompts often read like this: “Calculate the pH of a 0.100 M ethylamine solution if Kb = 5.6 × 10^-4” or “Calculate the pH of a 0.100 M ethylamine solution if pKb = 3.25.” The extra condition at the end tells you which equilibrium constant to use. If your homework truncates the wording, that is usually what is implied.

If the problem changes the concentration, the exact same method still works. Only the initial concentration changes in the ICE table. If the problem changes the temperature, then pKw may no longer equal exactly 14.00, and your pH conversion must use the given or appropriate pKw value.

Real-World Perspective

Ethylamine and related amines matter in industrial chemistry, pharmaceuticals, and synthesis. In solution, their basicity affects protonation state, reaction pathways, extraction behavior, and solubility. That is why pH calculations involving amines are more than classroom exercises. They are directly tied to practical chemistry decisions.

For trustworthy background on acid-base chemistry, water chemistry, and equilibrium concepts, see these authoritative educational resources:

• Chemistry LibreTexts for broad instructional chemistry coverage.
• U.S. Environmental Protection Agency (.gov) acid-base and alkalinity overview.
• Massachusetts Institute of Technology Chemistry (.edu) for university-level chemistry context.

Final Answer Summary

For a 0.100 M ethylamine solution at 25 degrees Celsius, using a representative Kb = 5.6 × 10^-4, the exact equilibrium calculation gives:

• [OH^–] ≈ 0.00722 M
• pOH ≈ 2.141
• pH ≈ 11.859
• Percent ionization ≈ 7.22%

If your course uses slightly different Kb or pKb values from another data table, your pH may differ by a few hundredths. That is normal. The method, however, remains exactly the same.