Solubility Practice Worksheet- Common Gases Exercises

What This Worksheet Actually Covers

You're here because you need to practice gas solubility problems and don't want to waste time on theory you'll never use. Fine. Let's get into it.

This worksheet focuses on Henry's Law applications for common gases you'll encounter in chemistry courses. If you're struggling with the basic concept, Henry's Law simply states that the amount of gas dissolved in a liquid is proportional to its partial pressure above that liquid. That's it. Everything else flows from there.

Common Gases and Their Solubility Data

Before you start crunching numbers, you need the right data. Here's what most textbooks use:

Notice CO₂ has a much higher constant than O₂ or N₂. That's why soda fizzes when you open the bottle—CO₂ escapes because it was under high pressure, and now it doesn't want to stay dissolved.

The Core Formula You Need

C = kP

Where:

• C = concentration of dissolved gas (mol/L)
• k = Henry's Law constant (specific to each gas and temperature)
• P = partial pressure of the gas (atm)

Some textbooks use C = kH × P where kH is the Henry's constant. Same thing, different symbol.

Practice Problems

Problem 1: Basic Oxygen Dissolution

Calculate the dissolved oxygen concentration in water at 298 K when the partial pressure of O₂ is 0.21 atm (normal atmospheric conditions).

Solution:

C = (1.3 × 10⁻³ M/atm) × (0.21 atm) = 2.73 × 10⁻⁴ mol/L

That's about 0.27 mM. Not much, but fish need it.

Problem 2: Carbonated Beverage

A soda bottle is pressurized at 4 atm with CO₂ at 298 K. What concentration of CO₂ dissolves in the liquid?

Solution:

C = (3.4 × 10⁻² M/atm) × (4 atm) = 0.136 mol/L

When you crack open the bottle, P drops to 0.0004 atm (ambient CO₂ partial pressure). The dissolved CO₂ comes out of solution. Physics doesn't care that you wanted to enjoy your drink.

Problem 3: Pressure Change Calculation

If water is saturated with N₂ at 5 atm and then the pressure is reduced to 1 atm, what happens to the dissolved nitrogen concentration?

Solution:

At 5 atm: C = (6.1 × 10⁻⁴) × 5 = 3.05 × 10⁻³ M

At 1 atm: C = (6.1 × 10⁻⁴) × 1 = 6.1 × 10⁻⁴ M

The excess nitrogen comes out of solution as bubbles. This is literally what happens in decompression sickness in divers. The bends aren't some mysterious affliction—it's Henry's Law in action.

How to Work Through These Problems

Stop guessing. Follow this sequence:

1. Identify what you're solving for. Usually concentration (C) or pressure (P).
2. Find the correct Henry's Law constant. Check the problem or table—don't guess from memory unless the constant is standard.
3. Plug in the values. Watch your units. If pressure is in kPa, convert to atm (1 atm = 101.325 kPa).
4. Calculate. Use scientific notation properly. 1.3 × 10⁻³ is correct; 1.3E-3 is fine on a calculator.
5. Check your answer. Does the magnitude make sense? If you get 500 M dissolved oxygen, something went wrong.

Common Mistakes to Avoid

• Using the wrong k value. Henry's constants vary wildly between gases. CO₂ is ~25 times more soluble than O₂ under the same pressure.
• Forgetting temperature dependence. These constants are for 298 K. Gas solubility drops as temperature rises—this is why warm water holds less dissolved oxygen.
• Ignoring unit conversions. Partial pressure in mmHg instead of atm will give you the wrong answer every time.
• Confusing mole fraction with molarity. Some k values are defined differently. Check which definition your textbook uses.

Why Temperature Matters

Every chemistry student learns that gas solubility decreases with increasing temperature. But do you know why?

Dissolving gas in liquid is an exothermic process. Heat is released when gas molecules interact with water. When you increase temperature, Le Chatelier's principle says the system shifts to counteract the change—by releasing gas. So hot water = less dissolved gas = fish die in warm discharge water from power plants.

This is also why you shouldn't boil water if you're trying to remove dissolved gases for an experiment. Let it equilibrate with the atmosphere instead, or bubble an inert gas through it.

Units Got You Confused?

Henry's Law constants appear in different forms depending on your source:

The dimensionless and atm forms are just inverses of the molarity form. If your constant is 769 atm, divide 1 by it to get 1.3 × 10⁻³ M/atm. Don't let the format trick you.

Real-World Context

Gas solubility isn't just exam fodder. It shows up in:

• Aquarium keeping — aeration increases O₂ partial pressure, driving more dissolution
• Carbonated drinks — CO₂ solubility under pressure keeps the fizz
• Environmental chemistry — ocean CO₂ absorption affects pH (ocean acidification)
• Medical hyperbaric chambers — high O₂ pressure dissolves more oxygen in blood
• Industrial degassing — removing dissolved gases from boiler feedwater prevents corrosion

Understanding the math lets you predict behavior. That's the point.

Download Your Worksheet

If you need structured practice, create your own worksheet using these problem templates:

• Given P and k, find C
• Given C and k, find P
• Given initial and final pressure, find concentration ratio
• Compare solubility of two gases under identical conditions
• Calculate gas volume released when pressure drops (using ideal gas law combined with Henry's Law)

Work through at least 10 problems before your exam. Pattern recognition comes from practice, not reading.

Quick Reference Cheat Sheet

Bookmark this:

• Higher k = more soluble (CO₂, NH₃ have high k values)
• Higher P = more dissolved gas (directly proportional)
• Higher T = less dissolved gas (inverse relationship)
• Units matter — always verify before calculating

That's everything you need for gas solubility problems. Now go practice.