Real-world examples of molarity and molality you’ll actually use

Instead of starting with theory, let’s jump straight into concrete, real examples of molarity and molality that show up in real life and real labs.

Think about:

• The saline solution used in IV drips at a hospital.
• The sulfuric acid solution inside a car battery.
• The antifreeze solution circulating in a car engine.
• The salt content of seawater affecting marine life.

Each of these is a perfect example of how concentration units are used. Some are best described using molarity (moles of solute per liter of solution), others using molality (moles of solute per kilogram of solvent), and the differences matter when temperature changes or when you need very precise colligative property calculations.

Classic lab examples of molarity and molality (with full calculations)

Let’s start with a few of the best examples of molarity and molality you’d actually see in a general chemistry lab.

Example 1: Preparing 0.10 M NaCl solution for a titration

You’re asked to prepare 250 mL of a 0.10 M sodium chloride (NaCl) solution.

Molarity setup

• Target molarity: 0.10 mol/L
• Volume: 250 mL = 0.250 L
• Molar mass of NaCl ≈ 58.44 g/mol

Moles of NaCl needed:

\[ n = M \times V = 0.10\,\text{mol/L} \times 0.250\,\text{L} = 0.0250\,\text{mol} \]

Mass of NaCl:

\[ m = n \times M_r = 0.0250\,\text{mol} \times 58.44\,\text{g/mol} \approx 1.46\,\text{g} \]

So you weigh out about 1.46 g of NaCl, dissolve it in some distilled water, and make the total volume up to 250 mL in a volumetric flask. This is one of the simplest textbook examples of molarity, but it mirrors what is done in teaching and research labs every day.

Turning the same solution into a molality example

Suppose you actually used 250.0 g of water (close to 250 mL at room temperature) instead of measuring volume.

• Mass of water (solvent): 250.0 g = 0.2500 kg
• Moles of NaCl: 0.0250 mol (from above)

Molality (m):

\[ m = \frac{n_{\text{solute}}}{m_{\text{solvent}}(\text{kg})} = \frac{0.0250}{0.2500} = 0.100\,\text{mol/kg} \]

Here, the numerical values of molarity and molality are close, but not identical if you account for solution density and temperature changes. This is a good example of examples of molarity and molality where both can be computed from the same preparation.

Example 2: 0.9% saline — medical molarity in action

Hospitals use 0.9% (w/v) NaCl solution for IV drips. That means 0.9 g NaCl in every 100 mL of solution.

Imagine you want the molarity of this saline solution.

• 0.9 g NaCl per 100 mL → 9.0 g per 1.00 L
• Molar mass NaCl ≈ 58.44 g/mol

Moles of NaCl in 1.00 L:

\[ n = \frac{9.0\,\text{g}}{58.44\,\text{g/mol}} \approx 0.154\,\text{mol} \]

So the molarity is:

\[ M \approx 0.154\,\text{mol/L} \]

This is a real example of molarity used in medicine, designed to be close to the osmolarity of human blood. Organizations like the U.S. National Library of Medicine and NIH discuss isotonic saline in clinical contexts, and the underlying chemistry is exactly this type of molarity example.

If the density of the saline is close to 1.00 g/mL, you can also estimate molality by treating 1.00 L as roughly 1000 g of solution and subtracting the solute mass. This makes saline a nice teaching example of examples of molarity and molality side by side.

Temperature-sensitive examples: why molality can be better

Molarity depends on volume, and volume changes with temperature. Molality depends on mass, which doesn’t. That’s why some of the best examples of molality show up in situations where temperature swings are large.

Example 3: Antifreeze in a car radiator (ethylene glycol–water)

Ethylene glycol (C₂H₆O₂) is used as antifreeze. Suppose you mix:

• 620 g of ethylene glycol
• 1000 g of water

Moles of ethylene glycol

Molar mass of C₂H₆O₂ ≈ 62.07 g/mol

\[ n = \frac{620\,\text{g}}{62.07\,\text{g/mol}} \approx 9.99\,\text{mol} \]

Mass of water (solvent) = 1000 g = 1.000 kg

Molality:

\[ m = \frac{9.99\,\text{mol}}{1.000\,\text{kg}} \approx 9.99\,\text{mol/kg} \]

This is a textbook-style example of molality, but it’s grounded in real engineering: antifreeze must perform from below 0 °F on a winter morning to well above 200 °F in a hot engine. Engineers use molality in colligative property equations (freezing point depression and boiling point elevation) because it stays constant when the coolant expands or contracts.

For reference, you can explore colligative property discussions in general-chemistry materials from universities such as MIT OpenCourseWare or UC Berkeley Chemistry for more context.

Example 4: Ocean water and freezing point depression

Seawater is another strong example of examples of molarity and molality at work in nature.

Average ocean water has about 3.5% salt by mass, mostly NaCl. That’s roughly 35 g of salt per 1000 g of seawater.

If we approximate that 35 g is almost all NaCl and treat the remaining 965 g as water, we can estimate molality.

• Mass of NaCl: 35 g
• Molar mass NaCl ≈ 58.44 g/mol

Moles of NaCl:

\[ n \approx \frac{35}{58.44} \approx 0.599\,\text{mol} \]

Mass of water (solvent): 965 g = 0.965 kg

Molality:

\[ m \approx \frac{0.599}{0.965} \approx 0.62\,\text{mol/kg} \]

This molality helps explain why seawater freezes below 32 °F (0 °C). Again, molality is preferred here because the ocean’s volume changes with temperature and pressure, but the mass of water does not.

If you wanted a molarity example from the same system, you’d need density data for seawater (around 1.025 g/mL on average), then convert mass of solution to volume. This dual use makes seawater a neat example of examples of molarity and molality in environmental chemistry.

Industrial and energy-related examples of molarity and molality

Chemistry isn’t just beakers and pipettes. Here are examples include car batteries, industrial cleaners, and more.

Example 5: Car battery acid (sulfuric acid solution)

Lead–acid car batteries typically use sulfuric acid (H₂SO₄) around 4–5 M in concentration. Let’s work through a simplified molarity example.

Say you have a sulfuric acid solution that is 4.5 M. That means:

• 4.5 moles of H₂SO₄ per liter of solution.

If you want to know how many grams of H₂SO₄ are in 500 mL:

• Volume = 0.500 L
• Molar mass H₂SO₄ ≈ 98.08 g/mol

Moles of H₂SO₄:

\[ n = 4.5\,\text{mol/L} \times 0.500\,\text{L} = 2.25\,\text{mol} \]

Mass of H₂SO₄:

\[ m = 2.25\,\text{mol} \times 98.08\,\text{g/mol} \approx 221.\,\text{g} \]

This is a real example of molarity that matters for battery performance and safety. Battery manufacturers track concentration closely because it affects voltage and lifespan. Some advanced battery research, including newer chemistries, still relies on molarity-based descriptions of electrolytes.

If you know the density of the acid solution, you could convert this to molality as well, giving yet another example of examples of molarity and molality from the same physical system.

Example 6: Industrial cleaning solution (NaOH)

Sodium hydroxide (NaOH) is widely used in manufacturing, from paper pulping to soap production. Suppose a plant uses a 2.0 M NaOH solution.

Assume you’re asked to prepare 2.0 L of this solution.

• Target molarity: 2.0 mol/L
• Volume: 2.0 L
• Molar mass NaOH ≈ 40.00 g/mol

Moles of NaOH needed:

\[ n = 2.0\,\text{mol/L} \times 2.0\,\text{L} = 4.0\,\text{mol} \]

Mass of NaOH:

\[ m = 4.0\,\text{mol} \times 40.00\,\text{g/mol} = 160.\,\text{g} \]

If the process temperature fluctuates significantly, engineers might instead specify concentration in molality, especially when calculating boiling point elevation or vapor pressure changes in closed systems. That turns this industrial NaOH bath into another example of examples of molarity and molality in real engineering decisions.

Pharmaceutical and biomedical examples of molarity and molality

The health and biotech sectors are packed with real examples of molarity and molality, especially when solutions need to match body fluids or maintain stability.

Example 7: Glucose solution for IV nutrition

A common hospital solution is 5% (w/v) dextrose in water (5% D/W), used for IV fluids.

• 5 g glucose per 100 mL solution → 50 g per 1.00 L
• Molar mass of glucose (C₆H₁₂O₆) ≈ 180.16 g/mol

Moles of glucose in 1.00 L:

\[ n = \frac{50}{180.16} \approx 0.277\,\text{mol} \]

Molarity:

\[ M \approx 0.277\,\text{mol/L} \]

This is a straightforward example of molarity used in medicine. Clinical references such as Mayo Clinic and MedlinePlus discuss IV fluids and dextrose solutions; behind those protocols is exactly this type of molarity example.

If you assume the density is about 1.02 g/mL, you can estimate solution mass and convert to molality. That gives yet another paired example of examples of molarity and molality from one clinically important solution.

Example 8: Buffer solutions in biochemistry

Enzyme activity is extremely sensitive to pH, so biochemists use buffer solutions with carefully controlled molarity.

Imagine preparing a 0.20 M phosphate buffer at pH 7.4, similar to physiological conditions.

You might mix specific amounts of NaH₂PO₄ and Na₂HPO₄ so that the total phosphate concentration is 0.20 M. Here, molarity is the natural choice because reactions often depend on the volume of the solution and the molar concentration of each species.

Molality could be used, but in most biochemical labs, volumes and molarity dominate day-to-day work, especially when working with instruments calibrated in moles per liter. This makes buffer preparation one of the best examples of molarity in modern biochemistry and pharmaceutical research.

When to prefer molarity vs. molality (with real examples)

By now we’ve walked through several detailed examples of examples of molarity and molality. Let’s summarize when each one shines, using situations you might actually encounter.

• Use molarity when:

  • You’re working at or near room temperature and volume changes are minor.
  • You’re mixing solutions in volumetric flasks and measuring with pipettes.
  • You’re doing titrations, buffer preparations, or routine lab work.
  • Example of this: preparing 0.10 M NaCl, 0.9% saline (≈0.154 M), 2.0 M NaOH.

• Use molality when:

  • Temperature varies a lot (antifreeze, environmental systems, some industrial reactors).
  • You’re calculating colligative properties: freezing point, boiling point, vapor pressure.
  • You care about mass-based accuracy more than volume.
  • Examples include: ethylene glycol–water antifreeze (~10 m), seawater (~0.62 m), concentrated electrolyte solutions where density changes strongly with temperature.

Modern chemistry research and industry in 2024–2025 still lean heavily on both units. Electrolyte design for next-generation batteries, advanced desalination research, and climate-related ocean chemistry all rely on the same style of calculations you’ve seen in these examples.

FAQ: short answers built around real examples

What are some simple examples of molarity?

A few everyday lab examples of molarity include 0.10 M NaCl prepared in a volumetric flask, 0.9% saline IV solution (about 0.154 M NaCl), and a 2.0 M NaOH cleaning solution used in manufacturing. Each describes how many moles of solute are present per liter of solution.

Can you give an example of molality from real life?

Antifreeze in a car radiator is a classic real-life example of molality. A mixture of 620 g ethylene glycol with 1000 g water is about 10 mol/kg. This molality is used to estimate how much the freezing point and boiling point of the coolant will shift, which directly affects engine protection.

Why would a chemist choose molality instead of molarity?

Chemists choose molality when temperature changes are significant or when they’re working with colligative properties. Because molality is based on mass of solvent, it doesn’t change when the solution expands or contracts. That’s why seawater freezing, antifreeze performance, and some high-precision industrial processes are best examples of situations where molality is preferred.

Are there examples of solutions where both molarity and molality are useful?

Yes. Saline IV solutions, seawater, and battery electrolytes are strong examples of examples of molarity and molality being used together. Molarity is convenient for dosing and reaction stoichiometry, while molality is more reliable for temperature-dependent properties and thermodynamic calculations.

How do I recognize whether a problem is an example of molarity or molality?

Look for the cues in the wording. If the problem gives or asks for volume of solution, it’s almost always a molarity example. If it gives mass of solvent in kilograms, it’s a molality example. Questions about freezing point depression, boiling point elevation, or vapor pressure lowering are often built around molality, while titrations and buffer problems are usually examples of molarity.