Calculate Ph Of 35 M Ascorbic Acid Yahoo

Use this premium calculator to estimate the pH of concentrated ascorbic acid solutions, including the common query “calculate ph of 35 m ascorbic acid yahoo.” The tool uses the acid dissociation constant for vitamin C and solves the first dissociation equilibrium with a quadratic expression for a more chemically sound ideal-solution estimate.

Expert guide: how to calculate pH of 35 M ascorbic acid

If you are searching for “calculate ph of 35 m ascorbic acid yahoo,” you are probably trying to verify an answer from a forum, an old Q&A thread, or a homework discussion. The key chemistry idea is that ascorbic acid, better known as vitamin C, is a weak acid rather than a strong acid. That means it does not fully dissociate in water. Because of that, you do not simply set hydrogen ion concentration equal to 35 M and take the negative logarithm. Instead, you use the acid dissociation constant, usually expressed as pKa or Ka.

Ascorbic acid is diprotic, meaning it can donate two protons in two separate acid dissociation steps. In practice, the first dissociation dominates the pH in ordinary acidic solutions because the second dissociation is much weaker. The first pKa of ascorbic acid is commonly reported around 4.10 at room temperature, while the second pKa is around 11.6 to 11.8. For most introductory calculations, the first equilibrium is enough:

H₂A ⇌ H⁺ + HA^–

Ka₁ = [H⁺][HA^–] / [H₂A]

Step 1: convert pKa to Ka

The first thing to do is convert pKa to Ka using the relation:

Ka = 10^-pKa

For pKa = 4.10:

Ka ≈ 10^-4.10 ≈ 7.94 × 10^-5

Step 2: define the starting concentration

If the solution concentration is 35 M, we can call the initial acid concentration C = 35.0 mol/L. In an ICE table, suppose that x mol/L dissociates:

• Initial: [H₂A] = 35.0, [H⁺] = 0, [HA^–] = 0
• Change: [H₂A] = -x, [H⁺] = +x, [HA^–] = +x
• Equilibrium: [H₂A] = 35.0 – x, [H⁺] = x, [HA^–] = x

Step 3: write the equilibrium expression

Substitute the ICE values into the Ka expression:

Ka = x² / (35.0 – x)

This can be solved exactly with the quadratic equation:

x = (-Ka + √(Ka² + 4KaC)) / 2

Using Ka = 7.94 × 10^-5 and C = 35.0:

x ≈ 0.0527 M

Since x represents [H⁺] in this simplified first-dissociation model, the pH becomes:

pH = -log₁₀(0.0527) ≈ 1.28

Ideal weak-acid estimate for 35 M ascorbic acid: pH ≈ 1.28

Why the answer is not simply negative log of 35

A common error in forum posts is to treat every acid concentration as if it were a strong acid. If ascorbic acid were a fully dissociated monoprotic strong acid at 35 M, the naive estimate would imply [H⁺] = 35 M and pH = -log(35) ≈ -1.54. That is not the correct weak-acid equilibrium treatment. Ascorbic acid has a relatively modest Ka, so the fraction dissociated is very small compared with the formal concentration. In the ideal weak-acid model, only about 0.15% dissociates at 35 M.

Why 35 M is chemically unusual

While the equilibrium math above is the standard classroom method, there is an important real-world caveat: a 35 M ascorbic acid solution is extraordinarily concentrated and may not be physically realistic under ordinary conditions. At such high concentrations, several things happen:

1. Activity coefficients deviate strongly from ideal behavior.
2. The solvent is no longer behaving like dilute water.
3. Density and solution structure change significantly.
4. The formal concentration can exceed realistic solubility or practical preparation limits.

That means the computed pH of 1.28 should be presented as an idealized equilibrium estimate, not a guaranteed measured value. In practical chemistry, pH electrodes and activity corrections matter much more in concentrated systems than in dilute textbook examples.

Quick approximation versus quadratic solution

For weak acids, many students learn the shortcut:

[H⁺] ≈ √(KaC)

If we plug in the same values:

√(7.94 × 10^-5 × 35) ≈ 0.0527 M

That gives essentially the same pH, because x is still much smaller than C. For 35 M ascorbic acid, the approximation and exact quadratic answer are nearly identical in the ideal model. However, using the quadratic method is still better practice, especially when you want a calculator that can handle a wider range of inputs robustly.

Ascorbic acid acid-base data

Property	Typical value	Why it matters for pH
Molar mass of ascorbic acid	176.12 g/mol	Useful for converting between grams and molarity.
pKa1	About 4.10	Controls the main acidic behavior in ordinary aqueous solutions.
pKa2	About 11.6 to 11.8	Second proton is weakly released and usually negligible in acidic pH calculations.
Ka1	About 7.9 × 10^-5	Used directly in the equilibrium equation.
Fraction dissociated at 35 M, ideal model	Roughly 0.15%	Shows why [H^+] is far lower than the formal concentration.

Comparison with other acids

Context helps. Ascorbic acid is acidic enough to give a low pH, but it is still much weaker than classic strong acids such as hydrochloric acid. The table below compares first-acidity data for several common acids discussed in general chemistry.

Acid	Typical pKa or strength description	Interpretation
Hydrochloric acid	Strong acid	Essentially complete dissociation in dilute water.
Acetic acid	pKa ≈ 4.76	Weaker than ascorbic acid in the first dissociation step.
Ascorbic acid	pKa1 ≈ 4.10	Moderately weak acid; stronger first acidity than acetic acid.
Citric acid	pKa1 ≈ 3.13	Stronger first dissociation than ascorbic acid.

How to interpret the keyword phrase “calculate ph of 35 m ascorbic acid yahoo”

This search phrase usually reflects one of three scenarios. First, someone may be looking for a direct numerical answer to a chemistry assignment. Second, they may have seen contradictory answers online and want a cleaner explanation. Third, they may be confusing lowercase “m” for molality versus uppercase “M” for molarity. In many casual forum posts, users type “35 m” when they actually mean “35 M.” If strict unit notation is required, that distinction matters. This calculator assumes you mean molarity unless you choose a different unit.

What happens if the concentration is lower?

As concentration decreases, the pH rises because [H⁺] from weak-acid dissociation becomes smaller. At lower concentration, the ideal weak-acid model is also more reliable because non-ideal effects become less severe. That is why pH calculations in textbooks often use concentrations like 0.10 M or 0.010 M rather than values as extreme as 35 M.

Why the second dissociation is usually ignored here

The second acidity constant of ascorbic acid is tiny compared with the first. Once the first proton dissociates, the anion HA^– does not release the second proton appreciably in a strongly acidic environment. Because the first equilibrium already places the pH near 1 to 3 for many concentrated solutions, the second dissociation contributes very little additional hydrogen ion. For that reason, the first Ka controls the result for this calculator.

Best practice for chemistry students

• Write the acid equilibrium before plugging in numbers.
• Convert pKa to Ka carefully.
• Use an ICE table if you are unsure.
• Check whether the acid is strong or weak.
• Use the quadratic method when you want the safest general solution.
• State whether your answer is an ideal estimate or a measured pH expectation.

Authoritative references for acid-base data and pH concepts

For readers who want more than a forum answer, these sources are much better than random message boards:

• NIH PubChem entry for ascorbic acid
• LibreTexts Chemistry educational resource
• U.S. EPA explanation of pH
• Michigan State University acid-base reference material

Final takeaway

If you need a practical answer to the query “calculate ph of 35 m ascorbic acid yahoo,” the standard ideal weak-acid calculation gives a pH of about 1.28 using pKa1 ≈ 4.10. That result comes from weak-acid equilibrium, not from assuming complete dissociation. Still, because 35 M is an extremely concentrated and likely non-ideal system, the actual measured pH in a real prepared sample could differ from the textbook estimate. For educational purposes, though, pH ≈ 1.28 is the correct and defensible calculation.

This page is for educational estimation only. In concentrated real systems, measured pH can diverge from ideal-equilibrium predictions because pH depends on hydrogen ion activity, not just formal concentration.