Calculate the pH of 0.15 M HCl and 0.12 M HNO3
This calculator computes hydrogen ion concentration and pH for two common strong monoprotic acids: hydrochloric acid and nitric acid. For ideal introductory chemistry conditions, both acids dissociate essentially completely, so the pH comes directly from the molarity using pH = -log10[H+].
Results
Enter or keep the default concentrations and click Calculate to see the pH of 0.15 M HCl and 0.12 M HNO3.
Expert guide: how to calculate the pH of 0.15 M HCl and 0.12 M HNO3
If you need to calculate the pH of 0.15 M hydrochloric acid and 0.12 M nitric acid, the good news is that this is one of the most direct acid-base calculations in general chemistry. Both HCl and HNO3 are classified as strong acids in water. That means they dissociate essentially completely under standard classroom conditions, so the hydrogen ion concentration is taken to be equal to the stated molarity of the acid, assuming one acidic proton per molecule. Since each of these acids is monoprotic, each mole of acid contributes approximately one mole of H+.
The key idea is simple: once you know the hydrogen ion concentration, you can calculate pH with the standard logarithmic relationship pH = -log10[H+]. For 0.15 M HCl, [H+] is approximately 0.15 M. For 0.12 M HNO3, [H+] is approximately 0.12 M. From there, the computation is straightforward. The calculator above automates the math, but understanding the chemistry helps you check whether your answer is realistic and whether your assumptions are valid.
Step 1: Recognize that both acids are strong monoprotic acids
Hydrochloric acid and nitric acid are both treated as strong acids in aqueous solution. In most chemistry courses, this means their dissociation is effectively complete:
- HCl(aq) → H+(aq) + Cl-(aq)
- HNO3(aq) → H+(aq) + NO3-(aq)
The word monoprotic matters because it tells you each formula unit donates only one hydrogen ion. If you were working with sulfuric acid, for example, the stoichiometry would be more nuanced because H2SO4 can supply more than one proton. With HCl and HNO3, one mole of acid gives about one mole of hydrogen ions in the typical idealized treatment used in introductory calculations.
Step 2: Convert acid concentration to hydrogen ion concentration
Because both acids dissociate completely and each contributes one proton, the hydrogen ion concentration is taken as equal to the molarity:
- For 0.15 M HCl, [H+] = 0.15 M
- For 0.12 M HNO3, [H+] = 0.12 M
This is why strong acid pH problems are often easier than weak acid equilibrium problems. You do not usually need an ICE table, Ka expression, or quadratic formula. The only major mathematical step is taking the negative base-10 logarithm.
Step 3: Apply the pH formula
The pH definition is:
Now substitute each concentration.
- For 0.15 M HCl: pH = -log10(0.15) = 0.8239, which rounds to 0.824
- For 0.12 M HNO3: pH = -log10(0.12) = 0.9208, which rounds to 0.921
These values make sense chemically. Since both concentrations are well above 0.1 M and both are strong acids, the pH values should be well below 1. Also, the more concentrated acid should have the lower pH. Because 0.15 M is greater than 0.12 M, the HCl solution is expected to be more acidic than the HNO3 solution, and the results confirm that.
Final answers
- pH of 0.15 M HCl = 0.824
- pH of 0.12 M HNO3 = 0.921
Comparison table for the two solutions
| Solution | Formal concentration (M) | Assumed [H+] (M) | Calculated pH at 25 C | Relative acidity |
|---|---|---|---|---|
| Hydrochloric acid, HCl | 0.15 | 0.15 | 0.8239 | More acidic of the two |
| Nitric acid, HNO3 | 0.12 | 0.12 | 0.9208 | Less acidic than 0.15 M HCl |
Why the pH values are not identical
Students sometimes expect all strong acids to have the same pH, but that is only true when they have the same effective hydrogen ion concentration. Strength tells you how fully an acid dissociates. Concentration tells you how much acid is present. In this comparison, both substances are strong acids, but their concentrations are not the same. Since pH depends on [H+], the solution with the larger hydrogen ion concentration has the lower pH.
Another useful way to interpret the results is to compare proton concentrations directly. The ratio 0.15 / 0.12 is 1.25, meaning the HCl solution has 25 percent more hydrogen ions per liter under the ideal assumption. That is why its pH is lower. Because pH is logarithmic, a 25 percent increase in [H+] does not create a huge numerical pH difference, but it is still chemically significant.
Important assumptions behind the simple calculation
The clean textbook answer depends on a few assumptions. In many laboratory and classroom contexts these assumptions are perfectly acceptable, but advanced chemistry work sometimes requires corrections:
- Complete dissociation: HCl and HNO3 are treated as fully dissociated in water.
- Monoprotic behavior: Each acid contributes one mole of H+ per mole of acid.
- Ideal behavior: The calculation uses concentration rather than activity.
- Standard temperature: Most pH examples assume conditions near 25 C.
- Negligible water autoionization: The water contribution of 1.0 × 10^-7 M H+ is tiny compared with 0.12 M or 0.15 M.
In more advanced analytical chemistry, pH is formally defined in terms of hydrogen ion activity rather than simple molarity. At higher ionic strength, activity coefficients can shift the measured pH slightly away from the idealized textbook value. However, for a standard educational problem phrased as “calculate the pH of 0.15 M HCl and 0.12 M HNO3,” the expected answers are the values shown above.
Reference data for hydrochloric acid and nitric acid
| Property | Hydrochloric acid (HCl) | Nitric acid (HNO3) |
|---|---|---|
| Molar mass | 36.46 g/mol | 63.01 g/mol |
| Acid classification | Strong acid | Strong acid |
| Number of ionizable protons | 1 | 1 |
| Typical pKa value in water references | About -6.3 | About -1.4 |
| Main aqueous anion after dissociation | Cl- | NO3- |
| Calculated pH at the stated concentration | 0.8239 at 0.15 M | 0.9208 at 0.12 M |
Common mistakes to avoid
- Using the wrong sign in the logarithm. pH is the negative log of hydrogen ion concentration, not the positive log.
- Confusing acid strength with acid concentration. A strong acid can still have a higher pH than another strong acid if it is less concentrated.
- Forgetting that pH is logarithmic. A small numerical difference in pH can correspond to a meaningful difference in acidity.
- Assuming all acids are treated this way. Weak acids require equilibrium calculations and cannot generally be solved by direct substitution.
- Rounding too early. Keep extra digits during the log step, then round at the end.
How to check whether your answer is reasonable
A quick reasonableness check is often enough to catch calculator errors. Since both concentrations are between 0.1 and 1.0 M, the pH should lie between 0 and 1 for both solutions. Also, because 0.15 M is more concentrated than 0.12 M, the HCl solution should have a lower pH than the HNO3 solution. If you get a negative pH or a value above 1 for these exact concentrations in a basic textbook problem, you should recheck your logarithm entry and signs.
You can also estimate mentally. Since log10(0.1) = -1, the pH of a 0.1 M strong acid is 1. Concentrations above 0.1 M must therefore have pH values slightly below 1. That immediately tells you 0.15 M and 0.12 M should both produce answers just under 1, which matches 0.824 and 0.921.
Where to learn more from authoritative sources
If you want to verify chemical properties or read more about pH and acid behavior, these authoritative resources are useful:
Bottom line
To calculate the pH of 0.15 M HCl and 0.12 M HNO3, treat both as fully dissociated strong monoprotic acids. Set [H+] equal to the molarity, then apply pH = -log10[H+]. The resulting values are 0.824 for hydrochloric acid and 0.921 for nitric acid. The HCl solution is more acidic because it has the greater hydrogen ion concentration. For standard chemistry homework and many introductory lab calculations, that is the complete and correct method.
Note: Real experimental pH measurements can differ slightly from ideal calculated values because pH is formally based on hydrogen ion activity, not just molarity.