Calculate The Ph Of A 0.045 M Hbr Solution

Chemistry Calculator

Use this interactive calculator to find the pH, hydrogen ion concentration, hydronium concentration, hydroxide concentration, and pOH for a hydrobromic acid solution. Because HBr is a strong acid, it dissociates essentially completely in water under typical introductory chemistry conditions.

HBr pH Calculator

How to Calculate the pH of a 0.045 M HBr Solution

To calculate the pH of a 0.045 M HBr solution, the key idea is that hydrobromic acid, HBr, is treated as a strong acid in water. In general chemistry, a strong acid is assumed to dissociate completely. That means every mole of HBr contributes essentially one mole of hydrogen ions, often represented as H⁺ or more precisely H₃O⁺ in aqueous solution. Because of that behavior, the hydrogen ion concentration is approximately equal to the molar concentration of the acid itself.

HBr → H⁺ + Br^–
[H⁺] = 0.045 M
pH = -log₁₀([H⁺]) = -log₁₀(0.045) ≈ 1.35

So the pH of a 0.045 M HBr solution is approximately 1.35. That result makes sense chemically because the solution is strongly acidic. Any pH near 1 indicates a high concentration of hydrogen ions compared with neutral water, which at 25 degrees C has a pH of 7.

Step-by-Step Solution

1. Identify the acid: HBr is hydrobromic acid, a strong acid.
2. Assume complete dissociation in water.
3. Set the hydrogen ion concentration equal to the acid concentration: [H⁺] = 0.045 M.
4. Use the pH formula: pH = -log₁₀[H⁺].
5. Compute the logarithm: pH = -log₁₀(0.045) ≈ 1.3468.
6. Round appropriately: pH ≈ 1.35.

Quick answer: A 0.045 M HBr solution has a pH of about 1.35, a pOH of about 12.65, and an [OH^–] concentration of about 2.22 × 10^-13 M at 25 degrees C.

Why HBr Is Treated as a Strong Acid

Hydrobromic acid belongs to the common set of strong acids introduced in first-year chemistry: HCl, HBr, HI, HNO₃, HClO₄, HClO₃, and the first ionization of H₂SO₄. What matters for pH calculations is that the acid ionizes so extensively that the equilibrium lies overwhelmingly toward products. For practical classroom calculations, that means the concentration of undissociated HBr is negligible compared with the concentration of hydrogen ions produced.

This is why HBr is much easier to work with than a weak acid. For a weak acid such as acetic acid, you would need an acid dissociation constant, Ka, and usually an ICE table or approximation method. For HBr, none of that is necessary in standard problem solving. You simply match the molarity of HBr to the molarity of hydrogen ions.

The Fundamental Chemistry Behind the Answer

In water, free protons do not exist by themselves for any meaningful length of time. Instead, they associate with water molecules to form hydronium, H₃O⁺. That means the phrase hydrogen ion concentration is really shorthand for hydronium ion concentration. So in this solution:

• [H₃O⁺] ≈ 0.045 M
• [Br^–] ≈ 0.045 M
• pH = -log₁₀(0.045) ≈ 1.35

Since pH and pOH are related by pH + pOH = 14 at 25 degrees C, the pOH is:

pOH = 14.00 – 1.3468 = 12.6532

Then the hydroxide ion concentration comes from:

[OH^–] = 10^-12.6532 ≈ 2.22 × 10^-13 M

Comparison Table: HBr Concentration vs pH

One helpful way to understand your result is to compare it with other hydrobromic acid concentrations. Because pH is logarithmic, a small change in concentration does not produce a linear change in pH. Every tenfold change in hydrogen ion concentration shifts pH by 1 unit.

HBr Concentration (M)	Assumed [H^+] (M)	Calculated pH	Acid Strength Interpretation
1.0	1.0	0.00	Extremely acidic, highly concentrated strong acid
0.10	0.10	1.00	Strongly acidic
0.045	0.045	1.35	Strongly acidic, typical textbook calculation
0.010	0.010	2.00	Acidic, but ten times less concentrated than 0.10 M
0.0010	0.0010	3.00	Still acidic, lower hydrogen ion concentration

How Acidic Is pH 1.35 in Real Terms?

A pH of 1.35 is very acidic. Remember that the pH scale is logarithmic, not linear. A solution with pH 1 is ten times more acidic than a solution with pH 2 and one hundred times more acidic than a solution with pH 3, in terms of hydrogen ion concentration. Therefore, a 0.045 M HBr solution contains a substantial amount of hydronium ions and should be handled with appropriate laboratory precautions.

From a safety perspective, hydrobromic acid solutions can be corrosive. In a real lab, concentration, temperature, and handling procedures all matter. Eye protection, chemical-resistant gloves, and good ventilation are standard. The calculator here is educational and designed to help with chemistry computation, not replace a safety data sheet or laboratory protocol.

Comparison Table: pH Benchmarks and Typical Reference Points

Substance or Reference	Typical pH	Comparison to 0.045 M HBr
Battery acid range	About 0 to 1	0.045 M HBr is slightly less acidic than the strongest examples in this range but still very corrosive
Stomach acid	About 1.5 to 3.5	0.045 M HBr at pH 1.35 is at or below the lower end of typical gastric acidity
Lemon juice	About 2 to 3	0.045 M HBr is significantly more acidic
Pure water at 25 degrees C	7.0	Many orders of magnitude less acidic than HBr solution
Household ammonia	About 11 to 12	Strongly basic, opposite side of the pH scale

Common Mistakes When Solving This Problem

• Forgetting that HBr is strong. Students sometimes treat HBr like a weak acid and try to use Ka. For standard introductory calculations, that is unnecessary.
• Using the concentration directly as pH. pH is not 0.045. You must take the negative base-10 logarithm.
• Missing the negative sign. The formula is pH = -log₁₀[H⁺].
• Rounding too early. If you keep more digits until the end, you get pH = 1.3468, which rounds to 1.35.
• Ignoring units. The concentration must be in mol/L for the basic pH formula as used here.

What If the Problem Uses mM Instead of M?

If concentration is given in millimolar, convert to molar first. For example, 45 mM equals 0.045 M because:

45 mM = 45 × 10^-3 M = 0.045 M

After that conversion, the same strong-acid method applies and the pH is still about 1.35.

Why Water Autoionization Is Negligible Here

Pure water contributes only 1.0 × 10^-7 M hydronium ions at 25 degrees C. Compared with 0.045 M from HBr, that amount is tiny. The acid contribution is about 450,000 times larger than the hydronium concentration from water alone. That is why water autoionization can be ignored in this calculation. For very dilute strong acid solutions, especially near 10^-7 M, the contribution from water matters much more. But for 0.045 M HBr, it does not materially affect the answer.

Formula Summary for This Exact Question

1. Recognize HBr as a strong acid.
2. Set [H⁺] equal to the acid molarity.
3. Apply pH = -log₁₀[H⁺].
4. Round the result correctly.

[H⁺] = 0.045 M
pH = -log₁₀(0.045) = 1.3468
Final answer: pH ≈ 1.35

Authoritative Chemistry References

For additional background on acids, pH, and aqueous chemistry, see these reputable educational and government resources:

• LibreTexts Chemistry for broad college-level acid-base explanations.
• U.S. Environmental Protection Agency for pH fundamentals and environmental context.
• MIT Chemistry for academic chemistry resources and instruction.

Final Takeaway

If you need to calculate the pH of a 0.045 M HBr solution, the process is straightforward because HBr is a strong acid. The acid dissociates essentially completely, so the hydronium concentration is the same as the listed molarity. Taking the negative logarithm gives a pH of approximately 1.35. This value indicates a strongly acidic solution, far more acidic than common foods and many everyday acidic substances. The calculator above lets you verify that number instantly and also explore how changing concentration changes pH on a logarithmic scale.