Calculate The Ph Of Each Solution.1.55 10 2 M Hbr

Use this premium calculator to find pH, pOH, hydronium concentration, and bromide concentration for hydrobromic acid solutions. It is optimized for scientific notation problems such as 1.55 × 10^-2 M HBr.

HBr pH Calculator

For strong monoprotic acid HBr, we assume complete dissociation in water at typical problem concentrations, so [H₃O⁺] ≈ [HBr].

Expert Guide: How to Calculate the pH of 1.55 × 10^-2 M HBr

When a chemistry problem asks you to calculate the pH of a solution such as 1.55 × 10^-2 M HBr, the key idea is to identify the acid type first. Hydrobromic acid, HBr, is a strong acid. In introductory and most general chemistry settings, strong acids are treated as dissociating completely in water. That means every dissolved HBr formula unit contributes one hydronium ion to the solution. Because pH depends on hydronium concentration, the calculation becomes very direct.

The pH scale is logarithmic and is defined as pH = -log₁₀[H₃O⁺]. For a strong monoprotic acid like HBr, the hydronium concentration is approximately equal to the acid concentration, as long as the solution is not in an extremely dilute or highly nonideal regime. Therefore, if the concentration of HBr is 1.55 × 10^-2 M, then the hydronium concentration is also about 1.55 × 10^-2 M. Substituting that value into the pH equation gives pH = -log(1.55 × 10^-2) ≈ 1.81. That is the expected answer for this specific problem.

Step by Step Method

1. Identify HBr as a strong acid.
2. Recognize that HBr is monoprotic, meaning it donates one proton per molecule.
3. Set [H₃O⁺] = [HBr].
4. Convert the concentration to decimal form if needed: 1.55 × 10^-2 = 0.0155 M.
5. Use the pH formula: pH = -log(0.0155).
6. Evaluate to obtain pH ≈ 1.81.

This is one of the most common strong acid calculations in first-year chemistry because it reinforces the relationship between concentration, dissociation, and logarithms. It also shows why scientific notation is so useful: chemistry often deals with concentrations ranging from very large values down to extremely tiny numbers.

Why HBr Is Treated as a Strong Acid

Hydrobromic acid belongs to the standard list of strong acids typically memorized in chemistry courses. Strong acids dissociate essentially completely in aqueous solution. In practical problem solving, that means the equilibrium lies overwhelmingly toward ions rather than intact acid molecules. For HBr, the dissociation process can be written as:

HBr + H₂O → H₃O⁺ + Br^–

Because one HBr produces one hydronium ion, the stoichiometry is one-to-one. That is why a 0.0155 M HBr solution yields approximately 0.0155 M hydronium and 0.0155 M bromide.

Detailed Calculation for 1.55 × 10^-2 M HBr

Let us work the math carefully:

• Given concentration: [HBr] = 1.55 × 10^-2 M
• Because HBr is a strong acid: [H₃O⁺] = 1.55 × 10^-2 M
• pH = -log(1.55 × 10^-2)
• pH = 1.81 when rounded to two decimal places

You can also estimate this mentally. Since 10^-2 corresponds to pH 2 if the coefficient were exactly 1, and because 1.55 is greater than 1, the pH should be slightly less than 2. That quick check confirms that 1.81 is sensible.

Quantity	Value for 1.55 × 10^-2 M HBr	Meaning
HBr concentration	0.0155 M	Initial strong acid concentration
Hydronium concentration	0.0155 M	Equal to acid concentration for monoprotic strong acid
pH	1.81	Strongly acidic solution
pOH	12.19	Using pH + pOH = 14 at 25 degrees C
Bromide concentration	0.0155 M	Conjugate base produced during dissociation

Significant Figures and Rounding

In pH calculations, the number of decimal places in the pH is linked to the number of significant figures in the concentration. The concentration 1.55 × 10^-2 has three significant figures. Therefore, the pH is usually reported with three digits total but two decimal places, giving 1.81. This is standard laboratory reporting practice. If your instructor uses a slightly different convention, follow the course guideline, but in general 1.81 is the correct rounded answer.

Common Mistakes Students Make

• Forgetting the negative sign in the exponent. If you use 1.55 × 10² instead of 1.55 × 10^-2, you get a completely unrealistic result for an ordinary pH homework problem.
• Treating HBr like a weak acid. HBr is strong, so you do not need an ICE table for standard problems.
• Using natural log instead of log base 10. pH uses base-10 logarithms.
• Dropping the coefficient. 1.55 × 10^-2 is not the same as 1.00 × 10^-2; the coefficient changes the pH noticeably.
• Confusing pH and pOH. The pOH of this solution is 12.19, but the pH is 1.81.

How Strong Acids Compare

At equal molar concentrations, common monoprotic strong acids such as HCl, HBr, and HNO₃ produce nearly the same hydronium concentration in idealized general chemistry calculations. That means a 0.0155 M solution of each would have approximately the same pH. The identity of the conjugate anion differs, but the one-proton stoichiometry does not. The table below shows how pH changes with concentration for a generic monoprotic strong acid, which applies well to HBr under textbook conditions.

Strong Acid Concentration (M)	[H3O^+] (M)	Calculated pH	Relative Acidity
1.0 × 10^-1	0.100	1.00	Very acidic
1.55 × 10^-2	0.0155	1.81	Strongly acidic
1.0 × 10^-3	0.00100	3.00	Acidic
1.0 × 10^-5	0.0000100	5.00	Weakly acidic by pH scale, but still from a strong acid source

Real pH Reference Data

To put this answer in context, a pH of 1.81 is much more acidic than many everyday substances. Neutral water at 25 degrees C has pH 7.0. Typical black coffee often falls around pH 4.85 to 5.10, while acid rain is commonly discussed around pH 4.2 to 4.4, though values vary by region and event. A solution with pH 1.81 is therefore thousands of times more acidic than coffee in terms of hydronium concentration because the pH scale is logarithmic, not linear.

When the Simple Strong Acid Assumption Needs More Care

Although introductory chemistry problems use complete dissociation for HBr, advanced chemistry can consider activity, ionic strength, and deviations from ideal behavior at high concentrations. At extremely low concentrations, water autoionization may also matter. However, for a concentration like 1.55 × 10^-2 M, the standard classroom assumption is excellent. The acid concentration overwhelmingly determines hydronium concentration, so the straightforward method remains correct.

How This Calculator Handles the Problem

The calculator above is specifically designed for scientific notation entries like 1.55 × 10^-2. You enter the coefficient and exponent separately to avoid formatting mistakes. It then reconstructs the molarity, assumes complete dissociation for HBr, computes pH and pOH, and displays the concentrations of hydronium and bromide. The chart visualizes the relationship between the original acid concentration and the resulting ion concentrations, helping you see the one-to-one stoichiometric pattern.

Practical Formula Summary

• C = a × 10^b where a is the coefficient and b is the exponent
• For HBr: [H₃O⁺] = C
• pH = -log₁₀(C)
• pOH = 14.00 – pH at 25 degrees C
• [Br^–] = C

Authoritative Chemistry References

For high-quality supporting information on pH, acids, and aqueous chemistry, see these trusted educational and government resources:

• LibreTexts Chemistry for foundational acid-base explanations.
• U.S. Environmental Protection Agency (.gov) for pH context in environmental science.
• U.S. Geological Survey (.gov) for a clear overview of pH and water chemistry.
• Purdue University (.edu) for pH fundamentals.

Final Answer

If the intended problem is calculate the pH of 1.55 × 10^-2 M HBr, the correct result is:

pH = 1.81

That answer follows directly from the fact that HBr is a strong monoprotic acid, so the hydronium concentration equals the acid concentration for standard general chemistry calculations.