Calculate the pH When 59.0 mL of 0.229 M Hydrobromic Acid Is Given
Use this interactive acid calculator to determine moles of HBr, hydrogen ion concentration, and final pH. Hydrobromic acid is a strong acid, so it dissociates essentially completely in water under typical introductory chemistry assumptions.
Results
Enter values and click Calculate pH to see the full breakdown.
How to calculate the pH when 59.0 mL of 0.229 M hydrobromic acid is given
To calculate the pH when 59.0 mL of 0.229 M hydrobromic acid is provided, the most important concept is that hydrobromic acid, HBr, is a strong acid. In standard general chemistry, strong acids are treated as completely dissociated in water. That means each mole of HBr contributes essentially one mole of hydrogen ions, written as H+ or more accurately H3O+ in aqueous solution.
Because HBr is monoprotic and strong, the hydrogen ion concentration is equal to the acid molarity, as long as you are finding the pH of the original solution itself. In this problem, the concentration is already given as 0.229 M. Therefore:
[H+] = 0.229 M
pH = -log[H+] = -log(0.229) ≈ 0.64
So the pH of the original 0.229 M HBr solution is approximately 0.64. The 59.0 mL volume is useful if you also want to calculate the number of moles of acid present, but it does not change the pH of the undiluted solution because pH depends on concentration, not on total amount alone.
Step by step solution
- Convert the volume to liters if you want moles: 59.0 mL = 0.0590 L.
- Use moles = M × L: moles HBr = 0.229 × 0.0590 = 0.013511 mol.
- Because HBr is a strong acid, moles of H+ = moles of HBr = 0.013511 mol.
- For the original solution, [H+] = 0.229 M.
- Apply the pH formula: pH = -log(0.229) = 0.6402.
- Report the final answer with suitable significant figures: pH ≈ 0.640, often rounded to 0.64.
Why the volume appears in the problem
Students often wonder why the problem includes 59.0 mL if the pH comes directly from the molarity. The reason is that chemistry problems commonly provide volume because volume is needed for mole calculations, stoichiometry, titration work, dilution, and reaction planning. In this case, volume lets you find how many moles of HBr are physically present:
- Volume: 59.0 mL = 0.0590 L
- Molarity: 0.229 mol/L
- Moles of HBr: 0.229 × 0.0590 = 0.013511 mol
If the problem had asked for the pH after dilution or after mixing with another solution, then the volume would become crucial to determine the new hydrogen ion concentration. But for the pH of the original solution alone, the concentration already tells you what you need.
Hydrobromic acid as a strong acid
Hydrobromic acid belongs to the group of strong hydrohalic acids. In aqueous solution, it dissociates essentially completely:
HBr(aq) → H+(aq) + Br–(aq)
Since one HBr unit produces one hydrogen ion, the stoichiometric relationship is 1:1. This makes strong acid pH calculations much simpler than weak acid problems, where an equilibrium expression and an acid dissociation constant, Ka, must be used.
| Quantity | Value | Meaning in this problem |
|---|---|---|
| Volume | 59.0 mL = 0.0590 L | Used to compute total moles of HBr present |
| Molarity | 0.229 M | Equal to [H+] for strong monoprotic HBr |
| Moles of HBr | 0.013511 mol | Total amount of acid in 59.0 mL |
| pH | 0.6402 | Acidity of the original solution |
Formula review for pH calculations
The central equation is:
pH = -log[H+]
For a strong monoprotic acid like HBr:
[H+] = acid molarity
Therefore:
pH = -log(0.229) = 0.6402
You can also verify that this answer makes chemical sense. A solution with concentration in the tenths of a molar for a strong acid should have a pH below 1. That is exactly what we obtain here.
Common mistakes to avoid
- Using volume instead of concentration for pH: pH depends on [H+], not on mL alone.
- Forgetting to convert mL to L for moles: 59.0 mL must become 0.0590 L before multiplying by molarity.
- Treating HBr like a weak acid: in basic chemistry problems, HBr is modeled as fully dissociated.
- Using natural log instead of base-10 log: pH calculations use log base 10.
- Confusing pH with pOH: pH measures acidity from H+, not hydroxide concentration.
Comparison with other common strong acid concentrations
To better understand the result, it helps to compare 0.229 M HBr with other strong acid concentrations. Because pH changes logarithmically, even moderate concentration differences produce meaningful pH shifts.
| Strong acid concentration (M) | Approximate pH | Interpretation |
|---|---|---|
| 1.00 | 0.00 | Very concentrated strong acid solution |
| 0.229 | 0.64 | This HBr problem |
| 0.100 | 1.00 | Common benchmark in introductory chemistry |
| 0.0100 | 2.00 | Ten times less concentrated than 0.100 M |
| 0.00100 | 3.00 | Dilute but still distinctly acidic |
Real scientific context: acidity, pH scale, and measurement
The pH scale is logarithmic, which means every 1 unit change in pH corresponds to a tenfold change in hydrogen ion concentration. This is why a solution with pH 0.64 is substantially more acidic than one with pH 1.64. In practical laboratory work, such acidic solutions require chemical splash goggles, gloves compatible with corrosive reagents, and careful handling under proper lab protocols.
Educational and governmental references consistently explain that strong acids can produce very low pH values in water and that pH is tied to hydronium concentration. For further reading, you can consult:
- U.S. Environmental Protection Agency on pH
- Chemistry LibreTexts educational chemistry reference
- Princeton University acid-base overview
How moles connect to pH in this problem
The total moles present are often the bridge between simple pH calculations and more advanced problems. Here, the 59.0 mL sample contains 0.013511 mol HBr. If you were to dilute this exact amount into a larger final volume, the concentration would decrease, and the pH would rise. For example, if the entire amount were diluted to 1.00 L, then:
- Moles H+ = 0.013511 mol
- New [H+] = 0.013511 M
- New pH = -log(0.013511) ≈ 1.87
This illustrates a key lesson: moles stay the same during dilution, but concentration changes. As concentration decreases, the pH increases.
Strong acid versus weak acid calculations
It is useful to compare HBr with a weak acid such as acetic acid. If you were given 0.229 M acetic acid instead of 0.229 M HBr, you could not simply set [H+] equal to 0.229 M. Instead, you would need an equilibrium setup because only a fraction of the weak acid molecules ionize. This is exactly why strong acid problems are commonly introduced first in chemistry classes.
- Strong acid: complete dissociation, direct pH calculation
- Weak acid: partial dissociation, equilibrium required
- HBr: treated as strong in aqueous solution
- Result here: pH is found immediately from the molarity
Data perspective on pH and concentration
The pH scale is widely used in environmental monitoring, industrial chemistry, and education. Government resources such as the EPA use pH as a standard parameter for assessing water quality because it strongly influences corrosion, aquatic life, and chemical speciation. In basic lab chemistry, pH values below 2 are generally recognized as highly acidic, and 0.229 M HBr clearly falls into that category with a pH near 0.64.
| pH range | Relative acidity | Typical interpretation |
|---|---|---|
| 0 to 1 | Extremely acidic | Concentrated strong acids or similarly corrosive solutions |
| 2 to 4 | Strongly acidic | Acidic lab or industrial solutions |
| 7 | Neutral | Pure water at standard conditions |
| 10 to 12 | Strongly basic | Basic cleaning or laboratory solutions |
Final answer for this chemistry problem
If the question is simply, “calculate the pH when 59.0 mL of 0.229 M hydrobromic acid”, the answer is:
pH = -log(0.229) = 0.6402
Final pH ≈ 0.64
If you also need the amount of acid present:
- Moles HBr = 0.229 mol/L × 0.0590 L = 0.013511 mol
- Moles H+ = 0.013511 mol