Has this ever happened to you? You pull a bottle of soda out of the fridge, and once you open it, you barely hear the tell-tale "hissing" of the carbonation. When you take a sip, it's flat. As you try to enjoy your now non-carbonated beverage, you think to yourself, "What caused this?" The answer is gas solubility. In soda, the "fizz" or bubbles are carbon dioxide (CO2), which is what causes the carbonation. Keep reading to learn about gas solubility and to learn how soda becomes flat.
- This article is about gas solubility
- First, we will learn what gas solubility means and how gases dissolve in liquid
- Next, we will learn how temperature affects solubility based on the energy of the solution process
- Then, we will learn how pressure affects solubility and see a graph demonstrating this effect
- Lastly, we will summarize the trends in solubility we have learned
Gas Solubility In Liquids
Let's begin by looking at the definition of solubility.
Solubility is the ability for a substance (called the solute) to dissolve in a solvent. The resulting substance is called a solution
Gases can dissolve in liquids. However, their solubility depends on a few factors, which we will be discussing in detail.
So, how does this process work? Here is a simple diagram to explain:
Fig.1 A diagram of the solvation process
Basically, the solute particles are attracted to the solvent particles, and these interactions cause the solvation process to begin. The solute and solvent "expand" (i.e., the particles are moved apart) so that the solute "fits" into this solvent.
This process is usually exothermic, meaning that energy/heat is released. Here's what an energy diagram of this process looks like:
Fig.2 Gas dissolving in a liquid is usually an exothermic process
Basically, the heat energy (enthalpy (ΔH)) is higher in the individual solute and solvent, so when the solution is formed (which is lower in energy), heat is released. This helps explain how temperature affects solubility, which we will discuss in this next section.
Temperature Effect On Gas Solubility
When temperature increases, the solubility of gas decreases. When the temperature of a substance is raised, it gains kinetic energy, which is the energy of motion. This increase in energy allows the gas particles to overcome the attractive forces between it and the solvent, so it will escape the solution.
As I mentioned earlier, gas solvation is usually an exothermic process. Since the solution exists at a lower energy than the individual solute and solvent, increases in the total energy mean the solute/solvent state is favored over the solution state. Below is a graph showing the change in solubility for different gases in water:
Fig.3-Trends in gas solubility (in water) with temperature, licensed by CC 4.0,
But what about when the solvation process is endothermic (absorbs heat)? Usually, when we discuss gas solubility, we are referring to the solubility in water (like in the graph above), which is a polar solvent. However, when gases are dissolved in non-polar solvents, such as hexane (C6H14), the process is endothermic.
A graph comparing these two processes is shown below:
Fig.4-Difference in enthalpy change for exothermic and endothermic solvation
Basically, the energy released when the non-polar solution is formed (ΔH3) is less than the energy released when the solution has a polar solvent. This is because the solute-solvent interactions are weaker since polar species have stronger attractive forces (such as dipole-dipole or hydrogen bonding).
Because of this, adding heat (i.e., increasing the temperature) makes the non-polar solution more favorable because it exists at a higher energy level.
Pressure Effect on Gas Solubility
Now let's talk about pressure. The effect of pressure on gas solubility is dictated by Henry's law.
Henry's Law states that the amount of gas dissolved in a set volume of liquid (i.e. solubility) is directly proportional to the pressure of the gas.
In mathematical terms:
$$C_g=k_HP_g$$
Where Cg is the solubility of the gas, kH is the Henry's law constant, and Pg is the pressure of the gas.
There are two factors that cause this relationship: compression and equilibrium.
Let's start with compression.
When the pressure of a gas is increased, the gas particles are compressed. This gives more room for more gas particles to dissolve.
One way to think of it is like pressing down on clothes in a container or drawer to make more room for more clothes to be placed on top.
Our second factor has to do with the equilibrium of gas and liquid.
Initially, some gas particles will collide with the surface of the liquid and dissolve, while others will gain enough energy to leave the liquid. This causes an equilibrium where the concentration of gas particles in and out of the liquid does not change.
Fig.4-Initial dynamic equilibrium
Now let's look at what happens when the pressure is increased.
Fig.6-Increase in pressure causes a shift in equilibrium
When the pressure is increased, the concentration of the gas solute (i.e., non-dissolved gas) increases. This is because the number of gas particles is staying the same while the volume is decreasing.
Since the concentration of gas solute is increasing, the equilibrium shifts to counterbalance this change. Basically, more gas particles will dissolve so that the concentration/ratio of concentration is restored to equilibrium levels.
The opposite occurs when you open a can or bottle of soda. The can/bottle is highly pressurized so more carbon dioxide (CO2) could be dissolved. When you open it, the pressure is decreased since the volume is increasing. This is why your soda becomes flat over time, since the solubility of CO2 decreases as the pressure decreases, and CO2 escapes the solution making the soda flat.
Gas Solubility graph
Now that we understand how pressure affects solubility, let's look at a graph:
Fig.7-Pressure versus solubility graph
Depending on the identity of the gas, the change in pressure will be more or less significant.
If we look back at our Henry's law equation:
$$C_g=k_HP_g$$
The Henry's law constant (kH) is what accounts for this difference. It's essentially a "proportionality constant" that shows to what extent the pressure affects the solubility of different gases.
Trends Of Gas Solubility
In summary, let's look back at what factors affect the solubility of a gas:
Temperature
The higher the temperature, the more soluble a gas will be if the solvation process is exothermic (usually polar solvents)
The reverse is true (i.e., the lower the temperature, the more soluble the gas) when the solvation process is endothermic (usually non-polar solvents)
Pressure
An increase in pressure leads to an increase in solubility
The extent of this effect is different for each gas (shown by kH, Henry's constant)
Gas Solubility - Key takeaways
- Solubility is the ability for a substance (called the solute) to dissolve in a solvent. The resulting substance is called a solution
- Temperature affects solubility based on the energy of the solution process
- If the process is exothermic (releases heat), an increase in temperature decreases solubility
- This is usually for polar solvents
- If the process is endothermic (requires heat), an increase in temperature increases solubility
- This is usually for non-polar solvents
- Henry's Law states that the amount of gas dissolved in a set volume of liquid (i.e. solubility) is directly proportional to the pressure of the gas.
- In mathematical terms: $$C_g=k_HP_g$$ Where Cg is the solubility of the gas, kH is the Henry's law constant, and Pg is the pressure of the gas.
References
- Fig.3-Trends in gas solubility (in water) with temperature (https://commons.wikimedia.org/wiki/File:CNX_Chem_11_03_gasdissolv.png) by OpenStax licensed by CC BY 4.0 (https://creativecommons.org/licenses/by/4.0/deed.en)
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