Best examples of preparing solutions: practical concentration & dilution examples

Why real examples of preparing solutions matter

Textbook formulas for concentration and dilution are short. Real lab days are not. When you’re standing at the bench with a bottle of solid NaCl, a 1.0 M stock acid on the shelf, and a protocol that assumes you magically know what to do, you want examples of preparing solutions that look like your actual work.

In 2024–2025, more labs, classrooms, and even at‑home experimenters rely on reproducible protocols and digital lab notebooks. That means your calculations need to be:

• Transparent (someone else can follow them)
• Checkable (you can catch unit mistakes)
• Consistent with safety guidelines

The good news: once you master a few practical concentration & dilution examples, you can remix them for almost any scenario.

Core formulas hiding behind all these examples

Before we walk through the best examples of preparing solutions, let’s put the two workhorse relationships on the table:

• Molarity (M):
  \( M = \dfrac{\text{moles of solute}}{\text{liters of solution}} \)

• Dilution (same solute, different volume):
  \( M_1 V_1 = M_2 V_2 \)

Here \(M_1\) and \(V_1\) are the concentration and volume of your more concentrated solution (often the stock), and \(M_2\) and \(V_2\) are the values for the diluted solution.

Every one of the examples of preparing solutions: practical concentration & dilution examples below is just a variation on these ideas, plus careful unit handling.

Everyday lab molarity: examples of preparing solutions from solids

Let’s start with solid solutes, because this is where most students and new lab techs first meet concentration math.

Example 1: Preparing 250 mL of 0.10 M NaCl

You need 250 mL of 0.10 M sodium chloride for a general chemistry lab.

1. Convert volume to liters
   \(250\ \text{mL} = 0.250\ \text{L}\)

2. Calculate moles needed
   \(M = n/V \Rightarrow n = M \times V\)
   \(n = 0.10\ \text{mol/L} \times 0.250\ \text{L} = 0.0250\ \text{mol}\)

3. Convert moles to grams (molar mass NaCl ≈ 58.44 g/mol):
   \(m = n \times M_r = 0.0250 \times 58.44 \approx 1.46\ \text{g}\)

4. Bench steps
   Weigh about 1.46 g NaCl, transfer to a 250 mL volumetric flask, add distilled water to dissolve, then fill to the calibration line.

This is one of the simplest examples of preparing solutions from a solid: no dilution, just molarity.

Example 2: Preparing 1.0 L of 0.25 M glucose (C₆H₁₂O₆)

Glucose solutions are common in biology labs and medical simulations.

1. Volume: 1.0 L (already in liters).
2. Moles needed: \(n = 0.25\ \text{mol/L} \times 1.0\ \text{L} = 0.25\ \text{mol}\).
3. Molar mass of glucose ≈ 180.16 g/mol.
4. Mass: \(m = 0.25 \times 180.16 \approx 45.0\ \text{g}\).

Again, weigh ~45.0 g glucose, dissolve in ~800 mL water, then make up to exactly 1.0 L. This is a classic example of preparing solutions for biochemistry or physiology labs.

Using stock solutions: practical concentration & dilution examples

In real labs, you rarely prepare everything from scratch. You lean on stock solutions. That’s where the dilution equation becomes your best friend.

Example 3: Diluting 6.0 M HCl to make 500 mL of 1.0 M HCl

You have concentrated 6.0 M hydrochloric acid and need 500 mL of 1.0 M solution.

Use \(M_1 V_1 = M_2 V_2\):

• \(M_1 = 6.0\ \text{M}\) (stock)
• \(M_2 = 1.0\ \text{M}\) (target)
• \(V_2 = 500\ \text{mL} = 0.500\ \text{L}\)

Solve for \(V_1\):

[
V_1 = \frac{M_2 V_2}{M_1} = \frac{1.0 \times 0.500}{6.0} \approx 0.0833\ \text{L} = 83.3\ \text{mL}
]

Bench reality:

• Measure 83 mL of 6.0 M HCl into a volumetric flask.
• Add acid to water, never the other way around, to reach 500 mL total.

This is one of the best examples of preparing solutions safely: the math is simple, but the safety practice (acid into water) is non‑negotiable. The CDC and NIOSH both emphasize proper handling of corrosive chemicals in lab safety guidance (CDC/NIOSH).

Example 4: Making 50 mL of 0.20 M NaOH from a 2.0 M stock

Here’s a smaller‑scale practical concentration & dilution example you might see in a titration lab.

• \(M_1 = 2.0\ \text{M}\)
• \(M_2 = 0.20\ \text{M}\)
• \(V_2 = 50.0\ \text{mL} = 0.0500\ \text{L}\)

[
V_1 = \frac{M_2 V_2}{M_1} = \frac{0.20 \times 0.0500}{2.0} = 0.0050\ \text{L} = 5.0\ \text{mL}
]

So, pipette 5.0 mL of 2.0 M NaOH into a 50 mL volumetric flask and dilute to the mark. This is a clean example of using a concentrated stock to avoid weighing tiny masses of hygroscopic NaOH pellets.

Serial dilutions: examples include microbiology and analytical chemistry

When you need a range of concentrations—say for calibration curves, enzyme assays, or microbial plate counts—you don’t recalculate from scratch every time. You perform serial dilutions.

Example 5: Tenfold serial dilution series from 1.0 M to 0.0001 M

Imagine you’re preparing standards for an absorbance vs. concentration curve.

Start with 1.0 M stock. Prepare:

• Tube A: 0.10 M
• Tube B: 0.010 M
• Tube C: 0.0010 M
• Tube D: 0.00010 M

Each step is a 1:10 dilution.

How to do it in practice:

• Add 9.0 mL solvent to each of four labeled tubes.
• Add 1.0 mL of 1.0 M stock to Tube A → 10 mL of 0.10 M.
• Mix, then transfer 1.0 mL from Tube A to Tube B → 0.010 M.
• Repeat down the line.

This is one of the most widely used examples of preparing solutions: practical concentration & dilution examples in microbiology and clinical labs. Serial dilution math underpins many methods used in public health and medical testing, including some protocols described in NIH and CDC method documents (NIH, CDC).

Percent solutions: mass/volume and volume/volume examples

Not every protocol talks in molarity. In medicine, biology, and some industrial settings, you’ll see percent solutions.

Example 6: Preparing 100 mL of 5% (m/v) NaCl

A 5% (m/v) solution means 5 g of solute per 100 mL of solution.

For 100 mL of 5% (m/v) NaCl:

• Mass of NaCl = 5 g
• Final volume = 100 mL

Bench steps: weigh 5 g NaCl, dissolve in ~70 mL water, then adjust volume to 100 mL. This is a simple example of percent mass/volume solution prep, similar in concept to saline used in medical settings (though clinical normal saline is 0.9% w/v).

The Mayo Clinic and other medical sources often refer to IV fluids by percent concentration, such as 5% dextrose in water (Mayo Clinic). Understanding these examples of preparing solutions helps bridge the gap between chemistry class and healthcare practice.

Example 7: Preparing 70% (v/v) isopropyl alcohol for disinfection

Since COVID‑19, demand for hand sanitizers and surface disinfectants has stayed high. A common target is 70% isopropyl alcohol by volume.

Suppose you want 500 mL of 70% (v/v) isopropyl alcohol from 99% stock.

Definition: 70% (v/v) means 70 mL isopropyl alcohol per 100 mL solution.

Use the dilution idea:

• \(C_1 = 99\%\)
• \(C_2 = 70\%\)
• \(V_2 = 500\ \text{mL}\)

[
V_1 = \frac{C_2 V_2}{C_1} = \frac{70 \times 500}{99} \approx 354.5\ \text{mL}
]

So you’d measure about 355 mL of 99% isopropyl alcohol and add water until the total volume is 500 mL. This is a very real example of preparing solutions: practical concentration & dilution examples that people used worldwide during the pandemic and still use in 2025 for infection control.

For disinfection guidelines and concentration ranges, the CDC maintains updated recommendations on their site (CDC disinfection).

Very dilute solutions: ppm and ppb in environmental and health contexts

When concentrations are tiny—think lead in drinking water or trace contaminants—scientists use parts per million (ppm) or parts per billion (ppb).

For dilute aqueous solutions, 1 ppm ≈ 1 mg solute per liter of water.

Example 8: Preparing 1.0 ppm fluoride standard solution

You’re calibrating an ion‑selective electrode to measure fluoride in drinking water.

Goal: 1.0 ppm F⁻ in water.

Assuming density ≈ 1.0 g/mL, 1.0 ppm ≈ 1.0 mg/L.

To make 1.0 L of 1.0 ppm solution:

• Mass of fluoride ion needed ≈ 1.0 mg.

If your source is sodium fluoride (NaF, molar mass ≈ 41.99 g/mol, F⁻ ≈ 19.00 g/mol), then:

• Fraction of mass that is F⁻: 19.00 / 41.99 ≈ 0.452.
• Mass of NaF needed: \(1.0\ \text{mg F⁻} / 0.452 \approx 2.21\ \text{mg NaF}\).

Weigh ~2.2 mg NaF and dilute to 1.0 L. This kind of example of ppm preparation is standard in environmental labs following EPA and related guidelines.

Buffer solutions: real examples from biochemistry

Buffers are everywhere: blood, cell culture media, enzyme assays. They resist pH changes, and their preparation combines concentration, dilution, and acid–base chemistry.

Example 9: Preparing 0.10 M phosphate buffer at pH 7.00

A common buffer uses NaH₂PO₄ (acid form) and Na₂HPO₄ (base form). The Henderson–Hasselbalch equation guides the ratio, but once you know the ratio, the actual preparation becomes another practical concentration & dilution example.

Say you have 0.20 M stock solutions of both NaH₂PO₄ and Na₂HPO₄.

You want 1.0 L of 0.10 M total phosphate at pH 7.00.

• Total phosphate concentration: \([\text{acid}] + [\text{base}] = 0.10\ \text{M}\).
• Suppose the pH and pKa give you a ratio \([\text{base}]/[\text{acid}] ≈ 1\) (near equal amounts for this example).

So, approximately:

• \([\text{acid}] ≈ 0.050\ \text{M}\)
• \([\text{base}] ≈ 0.050\ \text{M}\)

From 0.20 M stocks for 1.0 L final volume:

[
V_\text{acid} = \frac{0.050 \times 1.0}{0.20} = 0.25\ \text{L} = 250\ \text{mL} ] [ V_\text{base} = 0.25\ \text{L} = 250\ \text{mL}
]

Add 250 mL of each stock, then add water to 1.0 L and check pH, adjusting slightly if needed.

Buffer calculations show up constantly in biochemistry and molecular biology labs, including protocols from universities like Harvard and other research institutions (Harvard University).

Common mistakes in concentration & dilution calculations

Even with all these examples of preparing solutions: practical concentration & dilution examples, a few errors keep showing up in lab notebooks:

• Forgetting to convert mL to L when using molarity
• Treating \(V_1\) as “volume of water to add” instead of “volume of stock solution”
• Ignoring significant figures and reporting unrealistically precise values
• Adding water to concentrated acid instead of acid to water (a real safety hazard)

The best examples of preparing solutions are the ones you can reconstruct from your notes a week later. Always write down:

• The formula you used
• The numbers and units you plugged in
• A short description of what you physically did at the bench

Pulling it together: patterns across all these real examples

Across these eight-plus examples of preparing solutions, a few patterns should stand out:

• You always start from a target: a concentration unit and a final volume.
• You convert that target into either moles (for molarity) or mass/volume relationships (for percent and ppm).
• You work backward to get either grams of solid or volume of stock.
• You finish with a physical protocol that respects safety and accuracy.

Once you see that the same logic works for NaCl, HCl, isopropyl alcohol, fluoride standards, and phosphate buffers, these practical concentration & dilution examples stop feeling like separate tricks and start looking like one coherent toolkit.

FAQ: Short answers built around real examples

Q1. What are some common examples of preparing solutions in a general chemistry lab?
Common examples of preparing solutions include making 0.10 M NaCl from solid salt, preparing 0.25 M glucose for osmotic pressure experiments, and diluting concentrated acids like 6.0 M HCl down to 1.0 M for titrations. These are standard training exercises because they use the same molarity and dilution logic you’ll reuse everywhere.

Q2. Can you give an example of using M₁V₁ = M₂V₂ correctly?
A clear example of using \(M_1 V_1 = M_2 V_2\) is preparing 50 mL of 0.20 M NaOH from a 2.0 M stock. You solve \(V_1 = (0.20 \times 0.0500)/2.0 = 0.0050\ \text{L}\), so you pipette 5.0 mL of stock and dilute to 50 mL total.

Q3. How do I avoid errors when doing dilution calculations?
Write the dilution equation, label each term, and check units. Make sure \(V_1\) is the volume of stock, not the volume of solvent. Compare your result with physical intuition: if you’re making a less concentrated solution, the volume of stock should be smaller than the final volume. Comparing new problems to familiar examples of preparing solutions: practical concentration & dilution examples is a good way to sanity‑check your work.

Q4. What is an example of a percent solution used in healthcare or biology?
A very common example is 5% (m/v) glucose or dextrose solutions used in IV fluids. To prepare 100 mL of 5% (m/v) glucose, you dissolve 5 g of glucose and bring the total volume to 100 mL. Another is 70% (v/v) isopropyl alcohol for disinfection, made by diluting higher‑purity alcohol with water.

Q5. Where can I find more real examples of solution preparation and lab safety guidance?
For more real‑world examples and context, check:

• CDC lab safety and disinfection pages for concentration‑related guidance
• NIH and university chemistry department resources for step‑by‑step solution prep examples
• Medical sites like Mayo Clinic for how concentration appears in IV fluids and medications

Using these sources alongside the examples of preparing solutions: practical concentration & dilution examples in this guide will give you both the math and the real‑life context to apply it confidently.