Calculate the pH of a 0.93 M Solution of C2H5NH3Cl
This premium acid-base calculator determines the pH of ethylammonium chloride solutions by treating C2H5NH3+ as the conjugate acid of the weak base ethylamine. Enter the concentration, choose your calculation method, and instantly view pH, pOH, hydronium concentration, and a comparison chart.
How to calculate the pH of a 0.93 M solution of C2H5NH3Cl
To calculate the pH of a 0.93 M solution of C2H5NH3Cl, you first identify what kind of compound you are dealing with. Ethylammonium chloride is a salt made from a weak base, ethylamine (C2H5NH2), and a strong acid, hydrochloric acid (HCl). The chloride ion does not appreciably affect pH because it is the conjugate base of a strong acid, but the ethylammonium ion, C2H5NH3+, does. That cation behaves as a weak acid in water, donating a proton to form hydronium ions. Because hydronium concentration controls acidity, the pH is less than 7.
The key chemistry is the acid dissociation equilibrium:
Since many students first meet this problem in general chemistry, it helps to remember a shortcut: when a salt contains the conjugate acid of a weak base, you can calculate its acid dissociation constant, Ka, from the base dissociation constant, Kb, of the weak base. The relationship is:
At 25 degrees Celsius, Kw is typically taken as 1.0 × 10-14. For ethylamine, a commonly used Kb value is 5.6 × 10-4. Substituting:
Now let the initial concentration of ethylammonium ion be 0.93 M. If x is the amount that ionizes, then at equilibrium:
- [C2H5NH3+] = 0.93 – x
- [C2H5NH2] = x
- [H3O+] = x
The equilibrium expression becomes:
Because Ka is very small, x is tiny relative to 0.93, so the approximation 0.93 – x ≈ 0.93 is excellent. Then:
That means the hydronium concentration is about 4.08 × 10-6 M. Therefore:
So the pH of a 0.93 M solution of C2H5NH3Cl is approximately 5.39 when you use Kb = 5.6 × 10-4 for ethylamine at 25 degrees Celsius. The exact quadratic solution gives virtually the same answer because the dissociation is so small.
Why C2H5NH3Cl is acidic in water
A common point of confusion is that salts are often imagined as neutral substances. In reality, salts can be acidic, basic, or neutral depending on the acid and base from which they were formed. Ethylammonium chloride comes from a weak base and a strong acid, so the cation hydrolyzes and lowers pH. This is an important classification rule in aqueous chemistry:
- Break the salt into ions: C2H5NH3+ and Cl–.
- Identify the source of each ion.
- Cl– is neutral because it is the conjugate base of strong acid HCl.
- C2H5NH3+ is acidic because it is the conjugate acid of weak base C2H5NH2.
- Therefore the overall solution is acidic.
This logic is broadly useful for ammonium and alkylammonium salts. If you know the parent amine is a weak base, then its protonated form will have a small but real Ka. In concentrated solutions such as 0.93 M, the pH may still be only moderately acidic, because the conjugate acid itself is weak. That is why the final pH lands near 5.4 rather than near 1 or 2.
Step by step method with the exact equation
The approximation method is usually sufficient, but it is good practice to know the exact solution. Start from:
Rearranging gives:
This is a quadratic equation in x. The physically meaningful root is:
Substituting Ka = 1.79 × 10-11 and C = 0.93 gives x ≈ 4.08 × 10-6 M, which matches the approximation to the displayed precision. The percent ionization is:
That tiny percent ionization confirms that the weak acid approximation is valid. In chemistry classes, the 5% rule is often used: if the estimated x is less than 5% of the initial concentration, the approximation is acceptable. Here the percentage is far below 5%, so the simpler square root approach is entirely justified.
Data table: key constants and calculated values
| Quantity | Symbol | Value used | Meaning in this problem |
|---|---|---|---|
| Salt concentration | C | 0.93 M | Initial concentration of C2H5NH3+ from complete salt dissociation |
| Base dissociation constant of ethylamine | Kb | 5.6 × 10-4 | Describes basicity of C2H5NH2 in water |
| Water ion product at 25 degrees Celsius | Kw | 1.0 × 10-14 | Used to convert Kb to Ka |
| Acid dissociation constant of ethylammonium | Ka | 1.79 × 10-11 | Describes acidity of C2H5NH3+ |
| Hydronium concentration | [H3O+] | 4.08 × 10-6 M | Determines pH directly |
| Calculated pH | pH | 5.39 | Final acidity of the 0.93 M solution |
Comparison with related nitrogen-containing conjugate acids
Comparing ethylammonium with other ammonium-type species gives helpful perspective. Stronger conjugate acids correspond to weaker parent bases, while weaker conjugate acids correspond to stronger parent bases. Ethylamine is a somewhat stronger base than ammonia, so ethylammonium is a somewhat weaker acid than ammonium. That trend is reflected in pKa values and in the pH of equally concentrated salt solutions.
| Conjugate acid | Parent base | Typical Kb of base | Calculated Ka of conjugate acid | Approximate pH at 0.93 M |
|---|---|---|---|---|
| NH4+ | NH3 | 1.8 × 10-5 | 5.6 × 10-10 | 4.14 |
| CH3NH3+ | CH3NH2 | 4.4 × 10-4 | 2.3 × 10-11 | 5.33 |
| C2H5NH3+ | C2H5NH2 | 5.6 × 10-4 | 1.79 × 10-11 | 5.39 |
| (CH3)2NH2+ | (CH3)2NH | 5.4 × 10-4 | 1.85 × 10-11 | 5.39 |
These values are typical textbook or reference values used in general chemistry calculations. Small differences can appear from source to source because dissociation constants are temperature dependent and may be reported with different significant figures. Still, the trend remains the same: alkylammonium ions are weak acids, and concentrated solutions often have pH values in the mildly acidic range.
Common mistakes when solving this type of pH problem
1. Treating the salt as a strong acid
C2H5NH3Cl is not hydrochloric acid. Although chloride comes from HCl, the salt itself only contributes acidity through the weak acid C2H5NH3+. That is why the pH is near 5.39 instead of close to 0.
2. Using Kb directly instead of converting to Ka
The tabulated constant often given in data tables is Kb for ethylamine, not Ka for ethylammonium. Since the reacting species in solution is C2H5NH3+, the acid equilibrium must be used. Always convert with Ka = Kw / Kb when needed.
3. Forgetting complete salt dissociation
Ionic salts are generally assumed to dissociate completely in introductory chemistry unless otherwise stated. Therefore a 0.93 M solution of C2H5NH3Cl gives an initial C2H5NH3+ concentration of 0.93 M.
4. Ignoring significant figures and scientific notation
Hydronium concentration values are very small, so scientific notation matters. If you round too early, pH can shift in the second decimal place. A good practice is to keep several extra digits in intermediate steps and round the final pH according to the data precision.
Practical interpretation of the result
A pH of about 5.39 means the solution is acidic, but not highly corrosive purely on the basis of pH. In laboratory settings, alkylammonium salts may still require careful handling because concentration, ionic strength, temperature, and the chemical nature of the solute can influence behavior beyond pH alone. From a pure acid-base perspective, though, the number tells you that the hydronium level is modest compared with strong acids. The solution is more acidic than pure water by roughly a factor of 400 in hydronium concentration, since neutral water at 25 degrees Celsius has [H3O+] = 1.0 × 10-7 M.
When the approximation method works best
The square root approximation x = √(KaC) is popular because it is fast. It works well when Ka is small and the initial concentration is not extremely dilute. In this problem, Ka is on the order of 10-11 and the concentration is nearly 1 M, so the approximation is excellent. If the solution were much more dilute, or if the acid were stronger, then the quadratic method would be safer. This calculator lets you choose either method, and for 0.93 M C2H5NH3Cl the two answers are effectively identical.
Authoritative references for acid-base constants and aqueous chemistry
For readers who want to verify equilibrium concepts and water chemistry fundamentals, these authoritative resources are useful:
- LibreTexts Chemistry for acid-base equilibrium explanations hosted by an educational institution network.
- NIST Chemistry WebBook for reliable chemical reference data from a U.S. government source.
- U.S. Environmental Protection Agency for background on pH measurement and aqueous chemistry relevance in environmental systems.
Final answer
Using Kb(C2H5NH2) = 5.6 × 10-4 and Kw = 1.0 × 10-14 at 25 degrees Celsius:
[H3O+] ≈ 4.08 × 10-6 M
pH ≈ 5.39
Therefore, the pH of a 0.93 M solution of C2H5NH3Cl is approximately 5.39.