Best real-world examples of equivalence point in acid-base titrations

Textbook definitions are forgettable. Real examples of equivalence point in acid-base titrations are what stick, because you can connect the numbers, the curve, and the indicator choice.

Think of the equivalence point as the “chemically balanced handshake” between acid and base: the exact moment when stoichiometry says, we’re even. It’s not always at pH 7, and that misunderstanding haunts a lot of students.

Below, we’ll walk through several concrete examples of equivalence point in acid-base titrations that you’d see in:

• General and AP Chemistry labs
• Analytical chemistry courses
• Water quality and environmental monitoring
• Pharmaceutical and food analysis

Classic example of equivalence point: HCl with NaOH (strong acid–strong base)

Let’s start with the cleanest case — the one most textbooks love.

You titrate 25.0 mL of 0.100 M HCl (strong monoprotic acid) with 0.100 M NaOH (strong base). The balanced equation is:

\[ \text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O} \]

Moles of HCl initially:

\[ n_{\text{HCl}} = 0.0250\,\text{L} \times 0.100\,\text{mol/L} = 2.50 \times 10^{-3}\,\text{mol} \]

To reach the equivalence point, you need the same moles of NaOH:

\[ V_{\text{NaOH}} = \frac{2.50 \times 10^{-3}\,\text{mol}}{0.100\,\text{mol/L}} = 0.0250\,\text{L} = 25.0\,\text{mL} \]

At the equivalence point, you’ve got a solution of NaCl in water. Both ions are spectators, so the solution is effectively neutral. The pH is very close to 7.0 at 25 °F—sorry, 25 °C (room temperature).

This is one of the best examples of equivalence point in acid-base titrations for showing why indicators like phenolphthalein or bromothymol blue work well: the pH jump is steep and passes through 7.

Weak acid–strong base: acetic acid with NaOH (equivalence point pH > 7)

Now let’s look at an example of equivalence point in acid-base titrations where the pH at equivalence is not 7.

You titrate 25.0 mL of 0.100 M acetic acid, CH₃COOH (a weak acid, \(K_a \approx 1.8 \times 10^{-5}\)), with 0.100 M NaOH.

Reaction:

\[ \text{CH}_3\text{COOH} + \text{OH}^- \rightarrow \text{CH}_3\text{COO}^- + \text{H}_2\text{O} \]

Stoichiometry to reach the equivalence point is identical to the HCl case: 25.0 mL of NaOH. But the chemistry after neutralization changes everything. At equivalence, the solution contains mainly sodium acetate, CH₃COONa.

The acetate ion is a weak base:

\[ \text{CH}_3\text{COO}^- + \text{H}_2\text{O} \rightleftharpoons \text{CH}_3\text{COOH} + \text{OH}^- \]

So the pH at the equivalence point is greater than 7 (often around 8.5–9 depending on concentrations). You’ve created a basic solution from a weak acid–strong base titration.

This is where indicator choice matters. You want an indicator whose color change range sits in the basic region, like phenolphthalein (pH ~8.2–10). This example of equivalence point in acid-base titrations is a favorite in analytical chemistry because it shows how the conjugate base controls the pH.

For more background on weak acids and their conjugate bases, the LibreTexts chemistry project has a solid explanation: https://chem.libretexts.org

Weak base–strong acid: ammonia with HCl (equivalence point pH < 7)

Flip the roles: now the base is weak, and the acid is strong.

You titrate 25.0 mL of 0.100 M NH₃ (aqueous ammonia) with 0.100 M HCl.

Reaction:

\[ \text{NH}_3 + \text{H}^+ \rightarrow \text{NH}_4^+ \]

Again, stoichiometry gives 25.0 mL of HCl at the equivalence point. But now your solution is mostly ammonium chloride, NH₄Cl. The ammonium ion is a weak acid:

\[ \text{NH}_4^+ + \text{H}_2\text{O} \rightleftharpoons \text{NH}_3 + \text{H}_3\text{O}^+ \]

That means the equivalence point pH is less than 7, often around 5.5–6 depending on concentrations and temperature.

Phenolphthalein is a poor choice here because it changes in the basic region. You’d pick an indicator like methyl orange or bromocresol green, which changes color in the acidic range. This is one of the best examples of equivalence point in acid-base titrations for showing how weak bases drag the equivalence point into acidic territory.

The U.S. EPA’s water quality resources discuss ammonia and ammonium speciation in natural waters, which is directly related to this acid–base behavior: https://www.epa.gov/wqc

Polyprotic acid example: titration of sulfuric acid

So far, we’ve only looked at monoprotic acids. Real systems often involve polyprotic acids, where you can have more than one equivalence point.

Consider titrating 25.0 mL of 0.100 M H₂SO₄ with 0.100 M NaOH.

Stepwise reactions:

\[ \text{H}_2\text{SO}_4 + \text{OH}^- \rightarrow \text{HSO}_4^- + \text{H}_2\text{O} \]
\[ \text{HSO}_4^- + \text{OH}^- \rightarrow \text{SO}_4^{2-} + \text{H}_2\text{O} \]

Total moles of acidic protons:

\[ n_{\text{H}^+} = 2 \times 0.0250\,\text{L} \times 0.100\,\text{mol/L} = 5.00 \times 10^{-3}\,\text{mol} \]

Since NaOH is 0.100 M, you need 50.0 mL to neutralize both protons. On the titration curve, you can often see two equivalence points:

• First equivalence point: after 25.0 mL NaOH (conversion to HSO₄⁻)
• Second equivalence point: after 50.0 mL NaOH (conversion to SO₄²⁻)

This polyprotic case is a strong teaching example of equivalence point in acid-base titrations because it forces you to think in terms of steps and speciation, not just a single dramatic pH jump.

Environmental monitoring example: alkalinity titration of natural water

Not all titrations are clean lab solutions. Water treatment plants and environmental labs routinely run acid–base titrations to measure alkalinity — the water’s ability to neutralize acid. Here, the equivalence point tells you how much bicarbonate, carbonate, and sometimes hydroxide are present.

A common method:

• Sample: Natural water containing primarily HCO₃⁻ and CO₃²⁻
• Titrant: Strong acid like H₂SO₄

The titration curve can show:

• A first equivalence point near pH ~8.3 (conversion of CO₃²⁻ to HCO₃⁻)
• A second equivalence point near pH ~4.5 (conversion of HCO₃⁻ to H₂CO₃/CO₂)

These two equivalence points let analysts back-calculate carbonate and bicarbonate concentrations. This is a very real example of equivalence point in acid-base titrations used in environmental chemistry and water regulation.

The U.S. Geological Survey (USGS) and EPA both describe alkalinity titrations for field and lab work; see: https://water.usgs.gov/owq/FieldManual/Chapter6/section6.6/

Pharmaceutical example: assay of aspirin (acetylsalicylic acid)

Pharmaceutical labs use titrations to verify active ingredient content. Aspirin (acetylsalicylic acid) can be analyzed via back-titration or direct titration after hydrolysis.

A simplified classroom-style experiment:

• Dissolve an aspirin tablet in a known excess of standardized NaOH
• Allow complete hydrolysis and neutralization of the acid groups
• Titrate the remaining NaOH with standardized HCl

Here, the equivalence point in the back-titration of NaOH with HCl tells you how much NaOH did not react with aspirin. From that, you calculate how much NaOH did react, and thus the moles of aspirin.

This is an indirect but very practical example of equivalence point in acid-base titrations applied to real products. The FDA and USP (U.S. Pharmacopeia) methods often rely on this style of acid–base assay.

For a deeper dive into titration-based assays in pharmaceuticals, you can explore university analytical chemistry course notes, such as those from MIT OpenCourseWare: https://ocw.mit.edu

Food chemistry example: acidity of vinegar

Vinegar is basically acetic acid in water, usually labeled as “5% acidity” in grocery stores. That percentage is not a guess; it’s measured.

A common lab experiment:

• Sample: 10.0 mL of commercial vinegar
• Titrant: 0.500 M NaOH
• Indicator: Phenolphthalein

You titrate until the faint pink color persists — that’s your endpoint, chosen to coincide with the equivalence point for a weak acid–strong base system. Using the volume of NaOH at the equivalence point, you compute the acetic acid concentration and compare it to the label.

This is one of the best examples of equivalence point in acid-base titrations for connecting everyday life to classroom chemistry. The same principles scale up in food quality labs.

Why the equivalence point pH changes: strong vs weak, mono- vs polyprotic

Looking across all these examples of equivalence point in acid-base titrations, a few patterns emerge:

• Strong acid + strong base → pH ≈ 7 at equivalence
• Weak acid + strong base → pH > 7 at equivalence (basic salt formed)
• Weak base + strong acid → pH < 7 at equivalence (acidic salt formed)
• Polyprotic acids/bases → multiple equivalence points, each reflecting a stepwise reaction

The pH at the equivalence point is controlled by the salt in solution and its hydrolysis, not by the original acid or base directly. That’s why you can’t just assume “equivalence = pH 7.”

This is also why titration curves are so heavily emphasized in modern analytical chemistry. Automated titrators and software (common in 2024–2025 labs) don’t just look for a color change; they analyze the full pH vs. volume curve and calculate the equivalence point from the inflection, which is especially helpful for weak systems or polyprotic acids.

Indicators and instruments: choosing how to spot the equivalence point

In all these examples of equivalence point in acid-base titrations, there are two practical ways to detect the equivalence point:

• Indicators: Organic dyes that change color over a specific pH range
• pH meters / automatic titrators: Electronic measurement of pH as titrant is added

You match the indicator’s transition range to the expected pH at equivalence:

• Strong acid–strong base (HCl–NaOH): bromothymol blue (pH 6.0–7.6) or phenolphthalein (pH 8.2–10) both work because the jump is steep
• Weak acid–strong base (acetic acid–NaOH): phenolphthalein is better; its range aligns with the basic equivalence pH
• Weak base–strong acid (NH₃–HCl): methyl orange or bromocresol green align with the acidic equivalence pH

In many modern labs (especially in pharmaceutical and environmental testing), pH meters and automatic titrators are preferred for precision and for generating full titration curves that clearly show the equivalence point.

For pH meter calibration and good practice, NIST and many university chemistry departments provide detailed guidance; for example, see NIST’s pH standards information: https://www.nist.gov

FAQ: examples of equivalence point in acid-base titrations

Q1: Can you give a simple classroom example of equivalence point in acid-base titrations?
A very common classroom example is titrating 0.100 M HCl with 0.100 M NaOH. When equal moles have reacted (same volume of equal concentrations), you reach the equivalence point at about pH 7, and the indicator flips color sharply.

Q2: What are some real examples of equivalence point in acid-base titrations outside the classroom?
Real examples include alkalinity titrations for drinking water and wastewater, acidity testing of vinegar in food labs, and assay of acidic drugs like aspirin in pharmaceutical quality control. In each case, the equivalence point volume is used to calculate the concentration of the acid or base in the original sample.

Q3: Is the endpoint always the same as the equivalence point?
Not exactly. The equivalence point is defined by stoichiometry; the endpoint is where you observe a change (like indicator color). Good experimental design makes the endpoint occur very close to the equivalence point, but they are conceptually different.

Q4: Why does the pH at the equivalence point sometimes differ from 7?
Because the salt formed at the equivalence point can act as an acid or a base. For weak acid–strong base titrations, the conjugate base makes the solution basic (pH > 7). For weak base–strong acid titrations, the conjugate acid makes the solution acidic (pH < 7).

Q5: How many equivalence points can a titration have?
Monoprotic systems (like HCl or NaOH) have one equivalence point. Polyprotic acids (like H₂SO₄ or H₃PO₄) or polyprotic bases can have two or more equivalence points, each corresponding to neutralization of one acidic or basic proton step.

Q6: Which indicators are best for strong acid–strong base titrations?
Indicators whose transition ranges fall near neutral pH, such as bromothymol blue or phenolphthalein, are typically used. Because the pH jump is very steep around the equivalence point, there’s flexibility, and several different indicators can work well.

Q7: How do automated titrators find the equivalence point without an indicator?
They record pH after each tiny addition of titrant and construct a titration curve in real time. The software then finds the inflection point (where the slope is maximum) or uses derivative methods to pinpoint the equivalence point, even in complex or noisy data.