Real‑world examples of vapor pressure lowering in solutions examples

Everyday examples of vapor pressure lowering in solutions examples

Let’s start with situations you’ve probably seen before, even if you didn’t realize vapor pressure was involved.

Salty ocean water vs. pure water

Take pure water at 25 °F above freezing (so 57 °F, or about 14 °C). Its equilibrium vapor pressure at 25 °C is about 23.8 mmHg. When you dissolve salt (NaCl) in water, you create a solution where some of the surface sites that used to be water molecules are now occupied by ions. Fewer water molecules can escape into the gas phase, so the vapor pressure drops.

Consider seawater with an approximate molality of 0.6 mol/kg for NaCl. Because NaCl dissociates into two ions, the effective solute particle molality is closer to 1.2 m. Using Raoult’s law in its idealized form:

[
P_{solution} = X_{solvent} P^0_{solvent}
]

The mole fraction of water in seawater is slightly less than 1. That tiny reduction is enough to lower the vapor pressure by a few tenths of a percent. It’s not dramatic, but it’s measurable and directly tied to colligative properties like boiling point elevation and freezing point depression.

This is a classic example of vapor pressure lowering in solutions examples that also connects to why ocean water evaporates a bit more slowly than pure water at the same temperature.

Sugary soft drinks and fruit juice

Sugary beverages are another everyday example of vapor pressure lowering in solutions examples. A typical soda might contain around 10–12% sugar by mass. Those sugar molecules don’t evaporate, but they occupy space among water molecules and reduce the mole fraction of water.

The result: the vapor pressure of water above the soda is lower than above pure water. That slightly reduces evaporation, which is one reason open sugary drinks don’t go flat or dry out quite as fast as you might expect. From a Raoult’s law perspective, the sugar is a nonvolatile solute, so the only component contributing to vapor pressure is water, and its mole fraction is less than 1.

In high-sugar foods (jams, jellies, syrups), this same effect helps control water activity, which in turn influences microbial growth. Food scientists often talk more about water activity than vapor pressure, but the physics is the same: solutes lower vapor pressure and reduce the tendency of water to escape into the air or into microorganisms.

For more on water activity and food preservation, see resources from the U.S. National Library of Medicine: https://www.ncbi.nlm.nih.gov

Automotive and industrial examples include antifreeze and coolants

Ethylene glycol in car radiators

Car antifreeze is one of the best examples of vapor pressure lowering in solutions examples that shows up in engineering. Modern engine coolants are often around 50% ethylene glycol (C₂H₆O₂) by volume mixed with water.

Ethylene glycol is nonvolatile at engine operating temperatures. When it’s dissolved in water, it lowers the mole fraction of water and therefore the vapor pressure of the solution. That matters for at least three reasons:

• It raises the boiling point of the coolant (boiling point elevation).
• It reduces vapor lock and cavitation in the cooling system.
• It decreases coolant loss due to evaporation.

Suppose you prepare a solution that is roughly 40% ethylene glycol by mole fraction and 60% water. Under ideal behavior,

[
P_{solution} = X_{water} P^0_{water} = 0.60 \times P^0_{water}
]

If the vapor pressure of pure water at a given temperature is 100 kPa, the solution’s vapor pressure would be about 60 kPa. Real systems deviate from ideality, but the trend is the same: more glycol, lower vapor pressure.

This is a very practical example of vapor pressure lowering in solutions examples that mechanics and automotive engineers rely on without necessarily calling it Raoult’s law.

Industrial brines and evaporation ponds

In chemical manufacturing, concentrated salt solutions (brines) are used in everything from chlorine production to food processing. Highly concentrated brines have dramatically lower vapor pressures than pure water.

In large solar evaporation ponds used for salt production, engineers must account for the fact that as the brine becomes more concentrated, the rate of evaporation slows because the vapor pressure of water above the solution keeps dropping. This is a real-world example of vapor pressure lowering in solutions examples that directly affects design calculations for pond size, residence time, and expected salt yield.

For background on industrial brine use and solution thermodynamics, you can explore open course materials from MIT: https://ocw.mit.edu

Laboratory examples of vapor pressure lowering using Raoult’s law

Nonvolatile solute in a lab beaker: glucose in water

Imagine you dissolve 1.00 mol of glucose (C₆H₁₂O₆, nonvolatile) in 9.00 mol of water. The total moles in the solution are 10.00.

• Mole fraction of water:
  \(X_{water} = \frac{9.00}{10.00} = 0.90\)

• Vapor pressure of pure water at 25 °C: about 23.8 mmHg

Using Raoult’s law:

[
P_{solution} = X_{water} P^0_{water} = 0.90 \times 23.8 \text{ mmHg} \approx 21.4 \text{ mmHg}
]

The vapor pressure has been lowered by about 2.4 mmHg compared with pure water. This simple setup is a textbook example of vapor pressure lowering in solutions examples that also sets up boiling point elevation and freezing point depression calculations.

Binary liquid mixtures: ethanol and water

When both components are volatile, Raoult’s law is applied to each component. Take an ethanol–water mixture where each component contributes to the total vapor pressure:

[
P_{total} = X_{ethanol} P^0_{ethanol} + X_{water} P^0_{water}
]

Adding a nonvolatile solute (like sugar) to this mixture will lower the mole fractions of both volatile components. That reduces the total vapor pressure.

This shows up in distillation: sugary or salty aqueous mixtures boil at slightly higher temperatures because their vapor pressures are lower at a given temperature. Chemists routinely see this effect in fractional distillation columns when separating ethanol–water mixtures in the presence of dissolved solids.

The National Institute of Standards and Technology (NIST) maintains detailed vapor pressure data for pure components and mixtures: https://webbook.nist.gov

Biological and medical examples include body fluids and IV solutions

Electrolyte solutions in the human body

Blood plasma and intracellular fluids are not pure water; they’re packed with ions (Na⁺, K⁺, Cl⁻, HCO₃⁻) and small solutes like glucose and urea. These solutes lower the vapor pressure of water in body fluids relative to pure water.

Osmotic balance, which is vital for cell volume regulation, is directly tied to these colligative effects. While clinicians usually talk about osmolarity and osmolality rather than vapor pressure, it’s the same physical story: more solute particles → lower vapor pressure → higher osmotic pressure.

Isotonic saline (0.9% NaCl) is a nice example of vapor pressure lowering in solutions examples used in medicine. The dissolved salt slightly lowers the vapor pressure of water and matches the osmotic conditions of blood plasma, helping prevent red blood cells from swelling or shrinking during IV therapy.

You can find more on body fluid composition and osmolarity in medical references from the National Institutes of Health: https://www.nih.gov

Pharmaceutical formulations and humidity control

Solid drug products often include excipients (fillers, binders, stabilizers) that interact with water. When moisture is present as a thin film or in microenvironments, dissolved excipients lower the vapor pressure of water in that micro-layer. That influences:

• How fast tablets pick up or lose water in storage
• Stability of moisture-sensitive active ingredients
• Shelf life under different humidity conditions

Regulatory guidelines for pharmaceutical stability testing (like those followed by the FDA and ICH) explicitly control temperature and relative humidity, which are directly tied to vapor pressure phenomena.

Environmental and climate-related examples of vapor pressure lowering

Salt on icy roads and evaporation

When road crews spread salt on ice, the main story is freezing point depression, but vapor pressure lowering is part of the same colligative package. The salty meltwater that forms has a lower vapor pressure than pure water.

That means, at a given winter temperature, salty puddles evaporate slightly more slowly than pure water would. It’s a subtle effect compared to melting the ice in the first place, but it’s another example of vapor pressure lowering in solutions examples that connects everyday observations to the underlying thermodynamics.

Salty lakes and evaporation rates

Highly saline lakes, like the Great Salt Lake in Utah or the Dead Sea, provide large-scale real examples of vapor pressure lowering in solutions examples. Their brines can be so concentrated that the vapor pressure of water above the surface is significantly reduced compared with freshwater lakes at the same temperature.

This affects:

• Evaporation rates
• Local humidity near the lake
• Long-term water balance and shrinkage trends

As climate patterns shift and evaporation intensifies in some regions, understanding how salinity lowers vapor pressure becomes part of predicting lake level changes and ecosystem impacts.

Worked Raoult’s law example of vapor pressure lowering

To tie everything together, here’s a clean numerical example of vapor pressure lowering in solutions examples using Raoult’s law.

Suppose you have a solution made by dissolving 2.00 mol of a nonvolatile solute (like sucrose) in 18.00 mol of water.

• Total moles = 2.00 + 18.00 = 20.00 mol
• Mole fraction of water:
  \(X_{water} = \frac{18.00}{20.00} = 0.90\)

At 25 °C, the vapor pressure of pure water, \(P^0_{water}\), is about 23.8 mmHg.

Using Raoult’s law:

[
P_{solution} = X_{water} P^0_{water} = 0.90 \times 23.8 \text{ mmHg} \approx 21.4 \text{ mmHg}
]

Lowering of vapor pressure:

[
\Delta P = P^0_{water} - P_{solution} = 23.8 - 21.4 = 2.4 \text{ mmHg}
]

That 2.4 mmHg drop might sound small, but it’s enough to:

• Raise the boiling point slightly
• Lower the freezing point
• Reduce the rate of evaporation

Scale this up to industrial brines or antifreeze, and you’re looking at shifts in boiling and freezing behavior that are large enough to matter for safety, equipment design, and environmental impacts.

FAQ: examples of vapor pressure lowering in solutions examples

What are some simple examples of vapor pressure lowering in everyday life?

Common examples include salty ocean water, sugary drinks, jams and syrups, and windshield washer fluid or engine coolant in cars. In each case, dissolved nonvolatile solutes reduce the mole fraction of the volatile component (usually water), lowering the vapor pressure.

Can you give an example of vapor pressure lowering with a calculation?

Yes. Dissolve 1.00 mol of a nonvolatile solute in 9.00 mol of water at 25 °C. The mole fraction of water is 0.90, and the vapor pressure of pure water is about 23.8 mmHg. Raoult’s law predicts a solution vapor pressure of about 21.4 mmHg, a clear example of vapor pressure lowering caused by the solute.

How is vapor pressure lowering used in industry?

Industrial examples of vapor pressure lowering in solutions examples include concentrated brines in chemical plants, solar evaporation ponds for salt production, antifreeze and coolants in automotive and power systems, and solvent mixtures in distillation columns. Engineers exploit lower vapor pressures to control boiling points, reduce evaporation losses, and manage safety margins.

Are there biological examples of vapor pressure lowering?

Yes. Blood plasma, intracellular fluids, and IV saline are all real examples. Dissolved ions and small molecules lower the vapor pressure of water, which is reflected in osmotic pressure and water activity. This underpins cell volume regulation and safe design of medical solutions.

Does vapor pressure lowering always follow Raoult’s law exactly?

Not always. Raoult’s law describes ideal solutions. Real solutions—especially those with strong solute–solvent interactions or high concentrations—can deviate. Activity coefficients are used to correct for non-ideal behavior. But as a first approximation and for many dilute solutions, Raoult’s law gives very useful examples of vapor pressure lowering in solutions examples that match experiments reasonably well.