Calculate the pH of a 0.045 M HBr Solution
Use this premium chemistry calculator to solve pH, hydrogen ion concentration, pOH, and hydroxide ion concentration for hydrobromic acid solutions. HBr is treated as a strong acid that dissociates completely in water under standard general chemistry assumptions.
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How to calculate the pH of a 0.045 M HBr solution
If you searched for how to calculate the pH of a 0.045 M HBr solution, the problem is usually much simpler than it first appears. Hydrobromic acid, written as HBr, is a strong acid. In most general chemistry and introductory analytical chemistry settings, strong acids are assumed to dissociate essentially completely in water. That means each mole of HBr contributes one mole of hydrogen ions, more precisely hydronium in aqueous solution, but many textbook problems use H+ notation for simplicity.
Because HBr is a monoprotic strong acid, the relationship between acid concentration and hydrogen ion concentration is direct:
So if the solution concentration is 0.045 M, then under the standard strong-acid assumption:
Once you know hydrogen ion concentration, pH comes from the standard logarithmic formula:
Substitute the value:
Rounded suitably, the pH of a 0.045 M HBr solution is 1.35. If your class or instructor wants three decimal places, report 1.347. This is the standard answer expected in most homework systems and solution walkthroughs.
Step-by-step solution
- Identify HBr as a strong acid.
- Recognize that strong acids dissociate completely in typical general chemistry problems.
- Set hydrogen ion concentration equal to acid molarity: [H+] = 0.045 M.
- Apply the equation pH = -log10[H+].
- Compute: pH = -log10(0.045) = 1.347.
- Round according to the required significant figures or decimal places.
Why HBr is treated as a strong acid
Hydrobromic acid belongs to the standard list of strong acids commonly memorized in first-year chemistry: HCl, HBr, HI, HNO3, HClO4, and H2SO4 for its first dissociation. The defining classroom assumption is that these acids ionize nearly 100% in dilute aqueous solution. This matters because weak acids require an equilibrium setup with a Ka expression, while strong acids usually do not.
For HBr, that means you can skip an ICE table in routine pH problems. Instead of solving for x at equilibrium, you can move immediately from molarity to hydrogen ion concentration. This is one of the reasons this question appears often in homework sets: it tests whether you can identify a strong acid and apply the pH definition correctly.
Common student mistake
The most frequent mistake is forgetting that pH uses a negative logarithm. Another common error is plugging in 4.5 instead of 0.045. Since molarity is given directly in mol/L, you must use 0.045 in the logarithm. If you enter 4.5, you would get a negative pH, which would not match this moderate strong-acid concentration problem.
Worked example for 0.045 M HBr
Let us walk through the exact calculation in a clean, exam-ready format.
- Write the dissociation equation: HBr → H+ + Br–
- Identify stoichiometry: one mole of HBr produces one mole of H+
- Use the given molarity: [H+] = 0.045 M
- Apply the pH formula: pH = -log(0.045)
- Calculate: pH = 1.346787…
- Round: pH = 1.35 or 1.347 depending on instructions
At 25 degrees C, you can also find pOH using the relation:
Then hydroxide concentration follows from:
Comparison table: strong acid concentration vs pH
The table below shows how pH changes for several monoprotic strong acid concentrations. Since HBr behaves like a strong monoprotic acid in this context, its pH follows the same mathematical pattern as HCl or HI at equal concentration.
| Acid concentration (M) | [H+] assuming full dissociation (M) | Calculated pH | Interpretation |
|---|---|---|---|
| 1.0 | 1.0 | 0.000 | Very strong acidity in introductory scale terms |
| 0.10 | 0.10 | 1.000 | Common benchmark for strong-acid examples |
| 0.045 | 0.045 | 1.347 | Your target problem |
| 0.010 | 0.010 | 2.000 | Ten times less concentrated than 0.10 M |
| 0.0010 | 0.0010 | 3.000 | Still acidic, but much less concentrated |
How this compares with weak acid calculations
Students sometimes overcomplicate HBr problems because they remember that many acid-base problems require a Ka expression, an ICE table, or quadratic equation. That is true for weak acids like acetic acid or hydrofluoric acid, but not for a standard strong-acid HBr problem in a general chemistry course.
| Feature | Strong acid like HBr | Weak acid like acetic acid |
|---|---|---|
| Dissociation assumption | Essentially complete | Partial equilibrium |
| Main setup | [H+] ≈ initial molarity | Need Ka and equilibrium expression |
| Need ICE table? | Usually no | Usually yes |
| Math difficulty | Direct logarithm | Often algebraic or quadratic |
| For 0.045 M solution | pH = 1.347 | Depends on Ka |
Significant figures and rounding rules
In pH calculations, decimal places in the pH are tied to the number of significant figures in the concentration. The concentration 0.045 has two significant figures, so many instructors would accept a pH reported to two decimal places, which is 1.35. If your software or teacher asks for greater precision, 1.347 is mathematically correct to three decimal places.
Best reporting choices
- Standard classroom answer: 1.35
- More precise calculator output: 1.347
- Hydrogen ion concentration: 4.5 × 10-2 M
Interpreting the result chemically
A pH of about 1.35 indicates a strongly acidic solution. This means the hydrogen ion concentration is high compared with neutral water, which at 25 degrees C has [H+] = 1.0 × 10-7 M and pH 7.00. The 0.045 M HBr solution is therefore much more acidic than everyday mildly acidic liquids such as black coffee or tomato juice.
You can understand this with a quick order-of-magnitude comparison. A pH of 1.35 is about 5.65 pH units below neutral. Since each pH unit corresponds to a factor of 10 in hydrogen ion concentration, this solution has around 105.65 times more hydrogen ions than neutral water. That is roughly 4.5 × 105, or about 450,000 times higher [H+] than neutral water under standard assumptions.
Real reference values and educational context
The pH scale is a standard concept in chemistry education and public science resources. At 25 degrees C, the pH scale is tied to the ionic product of water and the relation between hydrogen and hydroxide ion concentrations. Introductory chemistry courses commonly use the equation pH = -log[H+] for aqueous acid problems, especially when working with strong acids. Educational and government resources that discuss pH fundamentals, water chemistry, and acid-base behavior include the following authoritative references:
Although LibreTexts is not a .gov site, it is a widely used academic educational resource. The U.S. Geological Survey and the Environmental Protection Agency both provide trusted explanations of pH behavior in water systems, while college chemistry materials commonly reinforce the exact same formulas used in this calculator.
Frequently asked questions about 0.045 M HBr pH
Is the pH exactly 1.35?
The unrounded value is about 1.3468. If your course uses two decimal places, report 1.35. If it requests three decimal places, report 1.347.
Why can I set [H+] equal to 0.045 M?
Because HBr is treated as a strong monoprotic acid in aqueous solution, it dissociates essentially completely. One mole of HBr gives one mole of H+.
Would this be different for H2SO4?
Yes. Sulfuric acid can release two protons, but the second dissociation is not handled the same way in all contexts. That makes it more nuanced than HBr in many course problems.
Do I need to consider water autoionization here?
No, not for a 0.045 M strong acid solution. The hydrogen ion contribution from water itself is negligible compared with 0.045 M.
Final answer
For a 0.045 M HBr solution, assuming complete dissociation as a strong acid:
That is the correct general chemistry result. Use the calculator above if you want an instant breakdown of pH, pOH, [H+], and [OH–] along with a visual chart.