Calculate the pH of a 0.00234 M HBr Solution
Use this interactive strong acid calculator to compute pH, hydrogen ion concentration, pOH, and hydroxide ion concentration for hydrobromic acid solutions.
HBr pH Calculator
Visual Analysis
This chart compares the hydrogen ion concentration, hydroxide ion concentration, and pH scale position for the entered HBr solution.
How to Calculate the pH of a 0.00234 M HBr Solution
To calculate the pH of a 0.00234 M HBr solution, you use one of the simplest acid-base relationships in general chemistry. Hydrobromic acid, written as HBr, is a strong acid. In dilute aqueous solution, strong acids are treated as fully dissociated, meaning essentially every dissolved HBr unit donates its proton to water. That makes the hydrogen ion concentration directly tied to the acid concentration. In practical classroom calculations, you assume that a 0.00234 M HBr solution has a hydrogen ion concentration, [H+], of 0.00234 M.
The pH definition is:
pH = -log10[H+]
Substitute the concentration:
pH = -log10(0.00234)
When evaluated, the result is approximately:
pH = 2.63
More precisely, the value is about 2.6308, but in many chemistry courses you report pH to two decimal places because the concentration has three significant figures. That gives a final answer of pH = 2.63. This confirms that the solution is distinctly acidic, as expected for a strong acid.
Why HBr Is Treated as a Strong Acid
Hydrobromic acid belongs to the family of hydrohalic acids. In water, HBr ionizes almost completely:
HBr + H2O → H3O+ + Br-
Because this reaction proceeds essentially to completion, HBr is grouped with strong acids like HCl, HI, HNO3, HClO4, and the first dissociation step of H2SO4. For pH calculations at ordinary instructional concentrations, the equilibrium expression is not usually needed. Instead, the concentration of the acid is taken to be the concentration of hydronium or hydrogen ions produced.
- Strong acid means nearly complete ionization in water.
- For monoprotic strong acids, one mole of acid gives one mole of H+.
- HBr is monoprotic, so 0.00234 M HBr gives approximately 0.00234 M H+.
- That direct relationship makes the pH calculation fast and reliable for standard coursework.
Step-by-Step Solution
Step 1: Identify the acid type
HBr is a strong monoprotic acid. The word monoprotic means each formula unit can donate one acidic proton. This matters because some acids, such as sulfuric acid, can donate more than one proton per molecule under some conditions. HBr contributes only one hydrogen ion per molecule.
Step 2: Write the dissociation assumption
Because HBr is strong, use:
[H+] = 0.00234 M
Step 3: Apply the pH formula
Use the definition:
pH = -log10[H+]
Step 4: Insert the concentration
pH = -log10(0.00234)
Step 5: Calculate
pH = 2.6308
Step 6: Report with appropriate precision
The concentration 0.00234 has three significant figures. In pH reporting, the number of decimal places should generally reflect the significant figures of the concentration. Therefore, the final reported pH is:
pH = 2.63
Related Quantities for a 0.00234 M HBr Solution
Once the pH is known, you can quickly derive other acid-base values. At 25 degrees C, the water ion-product relationship gives:
pH + pOH = 14.00
So:
pOH = 14.00 – 2.63 = 11.37
The hydroxide ion concentration is then:
[OH-] = 10^(-11.37) ≈ 4.27 × 10^-12 M
- Start with [H+] = 0.00234 M
- Compute pH = 2.63
- Find pOH = 11.37
- Compute [OH-] = 4.27 × 10^-12 M
| Property | Value for 0.00234 M HBr | Meaning |
|---|---|---|
| Acid concentration | 0.00234 M | Initial HBr concentration in solution |
| Hydrogen ion concentration, [H+] | 0.00234 M | Equal to acid concentration for strong monoprotic HBr |
| pH | 2.63 | Acidity level on the logarithmic pH scale |
| pOH | 11.37 | Complementary basicity measure at 25 degrees C |
| Hydroxide ion concentration, [OH-] | 4.27 × 10^-12 M | Very small because the solution is strongly acidic |
Comparison With Other Common Strong Acid Concentrations
Because pH is logarithmic, small concentration changes do not produce linear pH changes. Every tenfold increase in hydrogen ion concentration lowers pH by 1 unit. That logarithmic behavior is one of the most important ideas in acid-base chemistry. The table below shows how 0.00234 M HBr compares with several other common strong acid concentrations under the same complete dissociation assumption.
| Strong Acid Concentration (M) | [H+] (M) | Calculated pH | Relative Acidity |
|---|---|---|---|
| 0.100 | 0.100 | 1.00 | About 42.7 times more acidic than 0.00234 M in terms of [H+] |
| 0.0100 | 0.0100 | 2.00 | About 4.27 times more acidic than 0.00234 M |
| 0.00234 | 0.00234 | 2.63 | Reference case |
| 0.00100 | 0.00100 | 3.00 | Less acidic than 0.00234 M |
| 0.000100 | 0.000100 | 4.00 | Much less acidic than 0.00234 M |
Important Clarification About the Symbol M Versus m
In chemistry notation, uppercase M usually means molarity, or moles of solute per liter of solution. Lowercase m means molality, or moles of solute per kilogram of solvent. Many homework problems casually use the lowercase symbol when they really intend molarity, but technically these are different concentration units. If a problem literally says 0.00234 m HBr, the strict interpretation is molality.
For very dilute aqueous solutions, molality and molarity are often numerically close because water has a density near 1.00 g/mL. Under that approximation, a 0.00234 m HBr solution gives a pH very close to the result for 0.00234 M HBr, namely about 2.63. If you needed a high-precision value from molality, you would also need solution density to convert accurately to molarity before applying the pH formula.
Common Mistakes Students Make
Using the weak acid formula by accident
HBr is not a weak acid, so you do not set up a Ka expression for standard pH calculations. For a strong acid, the direct dissociation assumption is appropriate.
Forgetting the negative sign in the pH formula
The pH formula always contains a negative sign: pH = -log10[H+]. Since hydrogen ion concentrations in acidic solutions are often less than 1, their logarithms are negative. The minus sign converts the final pH into a positive number.
Mixing up significant figures and decimal places
When reporting pH, the number of decimal places should reflect the significant figures in the concentration. For 0.00234, there are three significant figures, so pH should typically be reported to three decimal places if carrying full precision in intermediate steps, or reasonably to two decimal places in standard coursework as 2.63.
Confusing pH with acid concentration directly
pH is not equal to concentration. Because pH is logarithmic, a concentration of 0.00234 M does not translate into a pH of 0.00234. Instead, it becomes the negative base-10 logarithm of the concentration.
Why the Result Makes Chemical Sense
A pH of 2.63 is definitely acidic but not as extreme as concentrated laboratory acid. For perspective, pure water at 25 degrees C has a pH near 7.00, while a 0.0100 M strong acid has a pH of 2.00. Since 0.00234 M is smaller than 0.0100 M, its pH should be higher than 2.00. The calculated value of 2.63 fits that expectation perfectly.
This type of reasonableness check is helpful in chemistry. Before finalizing any answer, ask whether the value should be acidic or basic, strong or weak, and larger or smaller than known benchmarks. Here, all those checks support the calculation.
When Water Autoionization Matters
In a solution as concentrated as 0.00234 M HBr, the contribution of hydrogen ions from water autoionization is negligible. Pure water contributes only about 1.0 × 10^-7 M hydrogen ions at 25 degrees C. That amount is tiny compared with 0.00234 M, so it does not meaningfully affect the calculation. This is another reason the direct strong-acid approximation works so well here.
Useful Reference Sources
If you want to verify pH concepts, acid strength classification, or chemical property data, these sources are helpful:
- U.S. Environmental Protection Agency: pH Overview
- National Institute of Standards and Technology Chemistry WebBook
- Purdue University General Chemistry Topic Review
Final Answer
For a standard general chemistry calculation, a 0.00234 M HBr solution is treated as a fully dissociated strong acid, so [H+] = 0.00234 M. Applying the pH equation gives:
pH = -log10(0.00234) = 2.63
Therefore, the pH of the solution is 2.63.