Calculate The Ph Of 0.0155 M Hbr

Chemistry Calculator

Use this premium interactive calculator to find the pH, hydronium concentration, pOH, and related acid strength details for hydrobromic acid solutions. The preset example is 0.0155 M HBr, a strong acid that dissociates essentially completely in water.

pH Calculator

How to Calculate the pH of 0.0155 M HBr

To calculate the pH of 0.0155 M HBr, the key idea is recognizing that hydrobromic acid is a strong acid. In standard general chemistry and analytical chemistry contexts, HBr dissociates essentially completely in water. That means every mole of HBr contributes approximately one mole of hydrogen ions, or more precisely hydronium ions, to the solution. Because HBr is monoprotic, the hydronium concentration is taken to be equal to the acid concentration for this type of classroom calculation.

For a 0.0155 M solution of HBr:

HBr(aq) → H⁺(aq) + Br^–(aq)

Since one mole of HBr produces one mole of H⁺, we set:

[H⁺] = 0.0155 M

Then apply the pH definition:

pH = -log[H⁺] = -log(0.0155) ≈ 1.81

So, the pH of 0.0155 M HBr is approximately 1.81. If your instructor wants more decimal places, the value is about 1.8097, which rounds to 1.810 to three decimal places. In most chemistry courses, writing pH = 1.81 is completely acceptable unless the problem specifically requests a particular rounding format.

Quick answer: 0.0155 M HBr is treated as a fully dissociated strong acid, so [H⁺] = 0.0155 M and pH = 1.81.

Step-by-Step Method

1. Identify the acid. HBr is hydrobromic acid, a strong acid in water.
2. Check proton count. HBr is monoprotic, meaning it releases one proton per formula unit.
3. Set hydronium concentration. Because dissociation is effectively complete, [H⁺] = 0.0155 M.
4. Use the pH equation. pH = -log(0.0155).
5. Round appropriately. pH ≈ 1.81.

Why HBr Is Treated as a Strong Acid

Students sometimes ask why they can directly substitute the molarity into the pH formula. The reason is that HBr belongs to the common list of strong acids taught in introductory chemistry. Strong acids ionize almost completely in dilute aqueous solution, so the equilibrium lies overwhelmingly toward products. This differs from weak acids like acetic acid or hydrofluoric acid, where a significant fraction remains undissociated and an equilibrium expression must be used.

For HBr, the bromide ion is the conjugate base, but it is extremely weak in water and does not meaningfully reverse the process under these conditions. As a result, the simple strong-acid approach is both chemically justified and educationally standard for concentrations like 0.0155 M.

Detailed Interpretation of the Result

A pH of 1.81 indicates a strongly acidic solution. Remember that the pH scale is logarithmic, not linear. A solution at pH 1.81 is far more acidic than one at pH 2.81, because a one-unit decrease in pH represents a tenfold increase in hydrogen ion concentration. This is why even concentrations that may appear numerically small, such as 0.0155 M, can still correspond to a very acidic solution.

At 25°C, the relationship between pH and pOH is:

pH + pOH = 14.00

Therefore, if pH ≈ 1.81, then:

pOH = 14.00 – 1.81 = 12.19

This large pOH value simply reflects the very low hydroxide concentration in a strongly acidic solution.

Common Mistakes When Solving This Problem

• Using an ICE table unnecessarily. For strong monoprotic acids such as HBr, complete dissociation is assumed in standard coursework.
• Forgetting the negative sign in the pH equation. pH is the negative logarithm of hydrogen ion concentration.
• Using natural log instead of base-10 log. In chemistry, pH uses log base 10.
• Misreading the concentration. 0.0155 M is not the same as 0.155 M or 0.00155 M.
• Confusing HBr with Br^–. The acid contributes H⁺, while bromide itself is not the acid source.
• Incorrect rounding. The calculator may display 1.8097, but depending on instructions you may report 1.81.

Comparison Table: Strong Acid Concentration and pH

Strong Acid Concentration (M)	Approximate [H^+] (M)	pH at 25°C	Interpretation
0.100	0.100	1.000	Very strongly acidic, often used in textbook examples.
0.0155	0.0155	1.810	The target HBr example in this calculator.
0.0100	0.0100	2.000	A familiar benchmark because log(10^-2) is easy to evaluate.
0.00100	0.00100	3.000	Still acidic, but ten times less acidic than pH 2.000.
0.000100	0.000100	4.000	Dilute acid solution where water autoionization may eventually become more relevant at even lower concentrations.

How 0.0155 M HBr Compares with Everyday pH References

Although laboratory acids should never be compared casually to household substances in terms of safe handling, pH references can help build intuition. Rainwater is typically near pH 5.6 in equilibrium with atmospheric carbon dioxide, pure water at 25°C is pH 7.0, and many acidic beverages fall around pH 2.5 to 3.5. A pH of 1.81 is substantially more acidic than those common examples, which is why proper laboratory safety matters.

Reference Solution or Condition	Typical pH	Relative Acidity vs 0.0155 M HBr	Notes
Pure water at 25°C	7.0	0.0155 M HBr is about 10^5.19 times higher in [H^+]	Neutral benchmark used in nearly all introductory chemistry courses.
Typical rainwater	5.6	0.0155 M HBr is about 10^3.79 times higher in [H^+]	Rain is naturally slightly acidic due to dissolved carbon dioxide.
Lemon juice	2.0 to 2.6	0.0155 M HBr is in a similar acidic range, often somewhat more acidic than the upper end	Food pH values vary by composition, dilution, and measurement method.
Household vinegar	2.4 to 3.4	0.0155 M HBr is more acidic in most cases	Vinegar contains acetic acid, which is a weak acid rather than a strong acid.

Significant Figures and Reporting

Chemistry instructors often care about how you report pH because the logarithm changes the way significant figures are handled. If the concentration is given as 0.0155 M, it has three significant figures. Therefore, the pH should typically be reported with three digits after the decimal if you are following strict logarithmic rules, giving 1.810. However, many textbook answers round to 1.81 for simplicity. If your course emphasizes formal reporting conventions, use 1.810; if not, 1.81 is usually accepted.

What If the Acid Were Weak Instead?

This example is easy because HBr is a strong acid. If the acid were weak, such as acetic acid, you could not simply equate molarity and [H⁺]. You would need the acid dissociation constant, K_a, write an equilibrium expression, and solve for the extent of ionization. That process often involves approximation or the quadratic formula. Recognizing whether an acid is strong or weak is therefore one of the most important first steps in any pH problem.

Classroom and Lab Context

In quantitative chemistry, pH calculations connect several core concepts: concentration, logarithms, equilibrium, and stoichiometry. For HBr in particular, the stoichiometric relationship is straightforward because it is monoprotic. More advanced coursework may discuss activity coefficients, nonideal behavior, and temperature-dependent values, but those refinements are typically unnecessary for general chemistry problems at this concentration. The standard educational result remains pH ≈ 1.81.

Reliable chemistry education and laboratory safety guidance can be found from authoritative public institutions. For broader scientific context, see the U.S. Geological Survey overview of pH at usgs.gov, Purdue University’s acid and base educational resources at chem.purdue.edu, and chemical safety information from the National Institute for Occupational Safety and Health at cdc.gov.

Worked Example Summary

1. Start with the given concentration: 0.0155 M HBr.
2. Recognize HBr as a strong monoprotic acid.
3. Set [H⁺] = 0.0155 M.
4. Compute pH = -log(0.0155) = 1.8097.
5. Round the result: pH = 1.81 or 1.810 depending on your class format.

Final Answer

If you need a concise final response for homework, quiz review, or lab preparation, you can write:

The pH of 0.0155 M HBr is approximately 1.81. Since HBr is a strong acid, it dissociates completely, so [H⁺] = 0.0155 M and pH = -log(0.0155) = 1.81.

This calculator automates that same logic while also showing hydronium concentration, pOH, and a concentration comparison chart. It is ideal for quickly checking strong-acid pH problems and reinforcing the concept that pH responds logarithmically to concentration changes.