Calculate The Ph Of 0.28 M Nh4Br Solution

Use this premium chemistry calculator to determine the pH of ammonium bromide solution from concentration and the base dissociation constant of ammonia. The tool shows the full acid-base logic, hydrogen ion concentration, pOH, and a visual chart of the species relationship.

NH4Br pH Calculator

Ammonium bromide is formed from a weak base (NH3) and a strong acid (HBr). Therefore, the solution is acidic because NH4+ hydrolyzes in water while Br- is essentially neutral.

How to Calculate the pH of 0.28 M NH4Br Solution

To calculate the pH of a 0.28 M NH4Br solution, you need to recognize what kind of salt ammonium bromide is. NH4Br is composed of the ammonium ion, NH4+, and the bromide ion, Br-. Bromide is the conjugate base of the strong acid HBr, so it does not significantly react with water. Ammonium, however, is the conjugate acid of the weak base NH3, so it does react with water to produce hydronium ions. That means an aqueous solution of NH4Br is acidic.

The key hydrolysis equilibrium is:

NH4+ + H2O ⇌ NH3 + H3O+

Because this reaction produces hydronium ions, the pH drops below 7. The problem is not solved by simply using the salt concentration as the hydrogen ion concentration. Instead, you first determine the acid dissociation constant of NH4+, then solve the equilibrium expression.

Step 1: Identify the Relevant Acid-Base Pair

NH4+ is the conjugate acid of NH3. Since the base dissociation constant of ammonia is usually known, we can calculate the acid dissociation constant of ammonium using the relationship:

Ka × Kb = Kw

At 25°C, typical textbook values are:

• Kb for NH3 = 1.8 × 10^-5
• Kw = 1.0 × 10^-14

So the acid constant for NH4+ is:

Ka = Kw / Kb = (1.0 × 10^-14) / (1.8 × 10^-5) = 5.56 × 10^-10

Step 2: Set Up the ICE Table

For a 0.28 M NH4Br solution, assume the initial ammonium concentration is 0.28 M. Let x be the amount of NH4+ that dissociates:

• Initial: [NH4+] = 0.28, [NH3] = 0, [H3O+] = 0
• Change: [NH4+] decreases by x, [NH3] increases by x, [H3O+] increases by x
• Equilibrium: [NH4+] = 0.28 – x, [NH3] = x, [H3O+] = x

Now apply the acid dissociation expression:

Ka = [NH3][H3O+] / [NH4+]

5.56 × 10^-10 = x^2 / (0.28 – x)

Step 3: Solve for Hydrogen Ion Concentration

Because Ka is very small, many instructors allow the weak acid approximation 0.28 – x ≈ 0.28. That gives:

x = √(Ka × C) = √((5.56 × 10^-10)(0.28))

x = 1.25 × 10^-5 M

Since x = [H3O+], the pH is:

pH = -log[H3O+] = -log(1.25 × 10^-5) ≈ 4.90

If you solve the quadratic equation exactly, the answer is essentially the same to normal reporting precision. So the pH of 0.28 M NH4Br solution is approximately 4.90.

Final Answer for 0.28 M NH4Br

Using standard 25°C constants, the calculated pH is about 4.90. This confirms that ammonium bromide solution is acidic, but not strongly acidic. The acidity arises entirely from the hydrolysis of NH4+ in water.

Why NH4Br Is Acidic and Not Neutral

Many students first learn that salts dissolve into ions and then assume every salt solution must be neutral. That idea works only for salts made from a strong acid and a strong base, such as NaCl. NH4Br is different because one ion is acid-active.

1. HBr is a strong acid, so Br- is a negligible base in water.
2. NH3 is a weak base, so its conjugate acid NH4+ has measurable acidity.
3. That weak acidity generates H3O+, lowering the pH below 7.

This is why NH4Br is classified as an acidic salt. The same logic applies to other ammonium salts such as NH4Cl and NH4NO3, assuming the anion comes from a strong acid and remains nonbasic.

Common Formulas Used in This Calculation

• Ka = Kw / Kb
• Ka = x^2 / (C – x)
• x ≈ √(KaC) for weak acids when x is small relative to C
• pH = -log[H3O+]
• pOH = 14 – pH at 25°C

Comparison Table: NH4Br pH at Different Concentrations

The table below uses the same standard constants, Kb(NH3) = 1.8 × 10^-5 and Kw = 1.0 × 10^-14, to show how pH changes as NH4Br concentration changes. These values are useful as reference points for labs, homework, and exam checking.

NH4Br Concentration (M)	Ka of NH4+	Approx. [H3O+] (M)	Approx. pH	Acid Strength Observation
0.010	5.56 × 10^-10	2.36 × 10^-6	5.63	Mildly acidic
0.050	5.56 × 10^-10	5.27 × 10^-6	5.28	Acidic
0.100	5.56 × 10^-10	7.45 × 10^-6	5.13	Acidic
0.280	5.56 × 10^-10	1.25 × 10^-5	4.90	Moderately acidic for a weak-acid salt
0.500	5.56 × 10^-10	1.67 × 10^-5	4.78	More acidic as concentration rises
1.000	5.56 × 10^-10	2.36 × 10^-5	4.63	Clearly acidic

Comparison Table: NH4Br Versus Other Salt Solutions

This second table helps place NH4Br in context. The values shown are representative 25°C expectations for similarly concentrated solutions when standard classroom assumptions are used.

Salt	Parent Acid	Parent Base	Expected Solution Character	Typical pH Trend Near 0.10 M
NaCl	HCl, strong acid	NaOH, strong base	Neutral	About 7.00
NH4Br	HBr, strong acid	NH3, weak base	Acidic	About 5.13
CH3COONa	CH3COOH, weak acid	NaOH, strong base	Basic	About 8.87
NH4CH3COO	CH3COOH, weak acid	NH3, weak base	Depends on Ka vs Kb	Often near neutral, but variable

Exact vs Approximate Method

For most general chemistry work, the approximation method is excellent for NH4Br because the percent ionization is tiny. In the 0.28 M case, the calculated hydronium concentration is around 1.25 × 10^-5 M, which is far smaller than 0.28 M. That means the assumption C – x ≈ C is justified.

Still, the exact method is more rigorous. Starting from:

Ka = x^2 / (C – x)

you rearrange to:

x^2 + Kax – KaC = 0

Then solve the quadratic equation for the positive root. The calculator above supports both methods, letting you compare them directly.

Typical Mistakes When Solving NH4Br pH Problems

• Treating NH4Br as a strong acid. The full salt concentration is not equal to [H3O+]. Only a tiny fraction hydrolyzes.
• Using Kb directly in the pH equation. The reacting species in solution is NH4+, so you need Ka for ammonium.
• Forgetting Br- is neutral. Bromide does not appreciably raise the pH.
• Skipping the log step. After finding [H3O+], you must convert to pH using the negative logarithm.
• Ignoring temperature dependence. If your Kb or Kw changes with temperature, your final pH changes too.

Lab and Exam Strategy

If you encounter a salt hydrolysis question under time pressure, classify the ions first. Ask whether the cation or anion comes from a weak acid or weak base. If the cation is the conjugate acid of a weak base, expect an acidic solution. If the anion is the conjugate base of a weak acid, expect a basic solution. If both come from strong parents, expect neutrality.

For NH4Br specifically, the fastest path is:

1. Recognize NH4+ is acidic.
2. Compute Ka = Kw / Kb.
3. Use [H3O+] ≈ √(KaC).
4. Take pH = -log[H3O+].

That workflow is reliable and efficient on quizzes, homework systems, and AP or college-level chemistry exams.

Authoritative References for Acid-Base Constants and Water Chemistry

U.S. Environmental Protection Agency: pH fundamentals
Chemistry LibreTexts: Acid-base properties of salts
University of Wisconsin: Salt hydrolysis tutorial

Bottom Line

To calculate the pH of 0.28 M NH4Br solution, convert the known Kb of ammonia into the Ka of ammonium, solve the weak acid equilibrium, and then convert hydronium concentration to pH. Using standard 25°C data, the result is pH ≈ 4.90. That makes NH4Br an acidic salt solution, not a neutral one. If you want a fast check, the calculator above gives both the exact and approximate result instantly and plots the concentration relationships visually.