Calculate pH of Solution 9.55 × 10-2 M HBr
Use this premium calculator to compute the pH, hydrogen ion concentration, pOH, and acidity classification for a hydrobromic acid solution. By default, it is set to the common chemistry problem: calculate pH of solution 9.55 × 10-2 M HBr.
Calculated Results
Enter or confirm the values above, then click Calculate pH.
How to Calculate pH of a 9.55 × 10-2 M HBr Solution
If you need to calculate the pH of a solution labeled 9.55 × 10-2 M HBr, the chemistry is straightforward because hydrobromic acid, HBr, is treated as a strong acid in introductory and general chemistry. That means it dissociates essentially completely in water:
HBr(aq) → H+(aq) + Br–(aq)Since one mole of HBr produces one mole of hydrogen ions, the hydrogen ion concentration is taken to be equal to the formal acid concentration. For the problem 9.55 × 10-2 M HBr, that gives:
[H+] = 9.55 × 10-2 M = 0.0955 MThe pH formula is:
pH = -log10[H+]Substitute the concentration into the equation:
pH = -log10(0.0955) ≈ 1.02Why HBr Is Treated as a Strong Acid
Hydrobromic acid belongs to the family of common strong acids that ionize almost completely in dilute aqueous solution. In most educational settings, HBr is grouped with HCl, HI, HNO3, HClO4, and the first ionization of H2SO4 as acids that do not require an equilibrium ICE table for basic pH calculations. Instead, you use direct stoichiometry and the negative logarithm of the resulting hydrogen ion concentration.
This matters because if the acid were weak, you would need a dissociation constant, often written as Ka, and the hydrogen ion concentration would have to be solved from an equilibrium expression. With HBr, that extra step is not necessary for standard problems. For a concentration such as 9.55 × 10-2 M, complete dissociation is an excellent approximation.
Step by Step Method for This Exact Problem
- Read the concentration correctly: 9.55 × 10-2 M means 0.0955 moles per liter.
- Identify HBr as a strong monoprotic acid.
- Set [H+] equal to the acid concentration: 0.0955 M.
- Apply the pH equation: pH = -log10(0.0955).
- Round appropriately based on significant figures: pH ≈ 1.02.
The logarithm is the only mathematical step that tends to trip students up. On a calculator, you should enter log(0.0955), which gives a negative number because the concentration is less than 1. Then apply the negative sign in front of the log result, yielding a positive pH near 1.02.
Common Student Mistakes When Solving 9.55 × 10-2 M HBr
- Using 9.55 as the concentration instead of 0.0955.
- Forgetting that 10-2 shifts the decimal two places left.
- Using pH = log[H+] instead of pH = -log[H+].
- Treating HBr like a weak acid and attempting an unnecessary equilibrium setup.
- Confusing pH with pOH.
If you avoid these errors, this calculation becomes one of the fastest acid-base problems you can solve. It is a classic example of how acid strength and concentration work together. Because the acid is strong, the concentration directly determines hydrogen ion availability. Because the solution is below 0.1 M but still substantial, the pH is low and clearly acidic.
Interpretation of the Result
A pH of about 1.02 indicates a very acidic solution. Remember that the pH scale is logarithmic, not linear. A difference of one pH unit corresponds to a tenfold change in hydrogen ion concentration. So a solution with pH 1 is far more acidic than a solution with pH 2. In practical chemical terms, a 0.0955 M HBr solution has a relatively high concentration of hydrogen ions and would be considered corrosive and hazardous in a laboratory environment.
| Solution Type | Approximate [H+] in mol/L | Approximate pH | Interpretation |
|---|---|---|---|
| Pure water at 25 degrees C | 1.0 × 10-7 | 7.00 | Neutral reference point |
| Rainwater | 2.5 × 10-6 | 5.60 | Slightly acidic due to dissolved carbon dioxide |
| Typical vinegar | 1.0 × 10-3 to 3.2 × 10-3 | 3.0 to 2.5 | Moderately acidic food acid system |
| 9.55 × 10-2 M HBr | 9.55 × 10-2 | 1.02 | Strongly acidic laboratory solution |
| 1.0 M strong acid | 1.0 | 0.00 | Extremely acidic idealized case |
Comparison with Other Strong Acids at the Same Concentration
Because HBr is a strong monoprotic acid, it behaves similarly to several other common strong acids at the same molarity. If each acid donates one proton per formula unit and fully dissociates, then the hydrogen ion concentration and pH are approximately identical for equal molar concentrations. This is why 0.0955 M HCl, HI, HNO3, and HBr all give nearly the same pH in standard textbook work.
| Acid | Acid Class | Molarity | Expected [H+] | Approximate pH |
|---|---|---|---|---|
| HBr | Strong monoprotic | 0.0955 M | 0.0955 M | 1.02 |
| HCl | Strong monoprotic | 0.0955 M | 0.0955 M | 1.02 |
| HI | Strong monoprotic | 0.0955 M | 0.0955 M | 1.02 |
| HNO3 | Strong monoprotic | 0.0955 M | 0.0955 M | 1.02 |
| H2SO4 | Strong first proton, more complex overall | 0.0955 M | Not handled by the same simple one proton assumption | Requires more careful treatment |
Significant Figures and pH Reporting
In chemistry, pH values are typically reported with decimal places that reflect the number of significant figures in the concentration. The concentration 9.55 × 10-2 has three significant figures, so the pH is commonly reported to two decimal places as 1.02. That presentation is both concise and consistent with standard chemistry lab formatting.
What About Temperature?
Most classroom pH calculations assume 25 degrees C, where pH + pOH = 14.00 and the ionic product of water is approximately 1.0 × 10-14. For a strong acid as concentrated as 0.0955 M, modest temperature changes do not materially alter the introductory result because the hydrogen ion concentration is dominated by the acid, not by water autoionization. The precise thermodynamic activity based pH can differ slightly from the idealized concentration based pH, but standard teaching problems almost always expect the simple answer 1.02.
Practical Chemistry Context
HBr is widely used in synthetic chemistry, inorganic chemistry, and industrial bromide preparation. In solution, it is highly acidic and reacts readily with bases. If you were to neutralize 0.0955 M HBr with an equal volume of 0.0955 M NaOH, the stoichiometric endpoint would occur when moles of H+ and OH– match. Before the endpoint, the solution stays acidic; after the endpoint, excess hydroxide determines the pH.
This calculator is especially useful for students checking homework, lab prework, online quizzes, and exam practice involving strong acids written in scientific notation. It converts the notation into decimal molarity, computes pH, derives pOH, and visualizes the relationship between acidity and hydrogen ion concentration.
Authoritative Learning Resources
For deeper study, consult authoritative chemistry resources from academic and government institutions:
- Chemistry LibreTexts educational resource
- United States Environmental Protection Agency for pH background and water chemistry context
- National Institute of Standards and Technology for chemical measurement standards
Quick Recap
- The notation 9.55 × 10-2 M equals 0.0955 M.
- HBr is a strong monoprotic acid, so [H+] = 0.0955 M.
- Use pH = -log10[H+].
- The result is pH ≈ 1.02.
- The solution is strongly acidic.
If your goal is simply to answer the problem “calculate pH of solution 9.55 10 2 m hbr,” the best concise answer is this: assume the concentration is 9.55 × 10-2 M HBr, convert to 0.0955 M, recognize HBr as a strong acid, and calculate pH = -log(0.0955) = 1.02. The interactive calculator above lets you verify that answer instantly and compare it visually with nearby concentration values.