What can speed up the dissolving process?
11.1 The Dissolution Process
An earlier chapter of this text introduced solutions, defined as homogeneous mixtures of two or more substances. Often, one component of a solution is present at a significantly greater concentration, in which case it is called the solvent. The other components of the solution present in relatively lesser concentrations are called solutes. Sugar is a covalent solid composed of sucrose molecules, C12H22O11. When this compound dissolves in water, its molecules become uniformly distributed among the molecules of water:
C 12 H 22 O 11 ( s ) ⟶ C 12 H 22 O 11 ( a q ) C 12 H 22 O 11 ( s ) ⟶ C 12 H 22 O 11 ( a q )
The subscript “aq” in the equation signifies that the sucrose molecules are solutes and are therefore individually dispersed throughout the aqueous solution (water is the solvent). Although sucrose molecules are heavier than water molecules, they remain dispersed throughout the solution; gravity does not cause them to “settle out” over time.
Potassium dichromate, K2Cr2O7, is an ionic compound composed of colorless potassium ions, K + , and orange dichromate ions, C r 2 O 7 2− . C r 2 O 7 2− . When a small amount of solid potassium dichromate is added to water, the compound dissolves and dissociates to yield potassium ions and dichromate ions uniformly distributed throughout the mixture (Figure 11.2), as indicated in this equation:
K 2 C r 2 O 7 ( s ) ⟶ 2 K + ( a q ) + C r 2 O 7 2− ( a q ) K 2 C r 2 O 7 ( s ) ⟶ 2 K + ( a q ) + C r 2 O 7 2− ( a q )
As with the mixture of sugar and water, this mixture is also an aqueous solution. Its solutes, potassium and dichromate ions, remain individually dispersed among the solvent (water) molecules.
Figure 11.2 When potassium dichromate (K2Cr2O7) is mixed with water, it forms a homogeneous orange solution. (credit: modification of work by Mark Ott)
Link to Learning
Visit this virtual lab to view simulations of the dissolution of common covalent and ionic substances (sugar and salt) in water.
Water is used so often as a solvent that the word solution has come to imply an aqueous solution to many people. However, almost any gas, liquid, or solid can act as a solvent. Many alloys are solid solutions of one metal dissolved in another; for example, US five-cent coins contain nickel dissolved in copper. Air is a gaseous solution, a homogeneous mixture of nitrogen, oxygen, and several other gases. Oxygen (a gas), alcohol (a liquid), and sugar (a solid) all dissolve in water (a liquid) to form liquid solutions. Table 11.1 gives examples of several different solutions and the phases of the solutes and solvents.
Different Types of Solutions
|soft drinks 1||CO2(g)||H2O(l)|
|hydrogen in palladium||H2(g)||Pd(s)|
|rubbing alcohol||H2O(l)||C3H8O(l) (2-propanol)|
Solutions exhibit these defining traits:
- They are homogeneous; after a solution is mixed, it has the same composition at all points throughout (its composition is uniform).
- The physical state of a solution—solid, liquid, or gas—is typically the same as that of the solvent, as demonstrated by the examples in Table 11.1.
- The components of a solution are dispersed on a molecular scale; they consist of a mixture of separated solute particles (molecules, atoms, and/or ions) each closely surrounded by solvent species.
- The dissolved solute in a solution will not settle out or separate from the solvent.
- The composition of a solution, or the concentrations of its components, can be varied continuously (within limits determined by the solubility of the components, discussed in detail later in this chapter).
The Formation of Solutions
The formation of a solution is an example of a spontaneous process , a process that occurs under specified conditions without the requirement of energy from some external source. Sometimes a mixture is stirred to speed up the dissolution process, but this is not necessary; a homogeneous solution will form eventually. The topic of spontaneity is critically important to the study of chemical thermodynamics and is treated more thoroughly in a later chapter of this text. For purposes of this chapter’s discussion, it will suffice to consider two criteria that favor, but do not guarantee, the spontaneous formation of a solution:
- a decrease in the internal energy of the system (an exothermic change, as discussed in the previous chapter on thermochemistry)
- an increased dispersal of matter in the system (which indicates an increase in the entropy of the system, as you will learn about in the later chapter on thermodynamics)
In the process of dissolution, an internal energy change often, but not always, occurs as heat is absorbed or evolved. An increase in matter dispersal always results when a solution forms from the uniform distribution of solute molecules throughout a solvent.
When the strengths of the intermolecular forces of attraction between solute and solvent species in a solution are no different than those present in the separated components, the solution is formed with no accompanying energy change. Such a solution is called an ideal solution . A mixture of ideal gases (or gases such as helium and argon, which closely approach ideal behavior) is an example of an ideal solution, since the entities comprising these gases experience no significant intermolecular attractions.
When containers of helium and argon are connected, the gases spontaneously mix due to diffusion and form a solution (Figure 11.3). The formation of this solution clearly involves an increase in matter dispersal, since the helium and argon atoms occupy a volume twice as large as that which each occupied before mixing.
Figure 11.3 Samples of helium and argon spontaneously mix to give a solution.
Ideal solutions may also form when structurally similar liquids are mixed. For example, mixtures of the alcohols methanol (CH3OH) and ethanol (C2H5OH) form ideal solutions, as do mixtures of the hydrocarbons pentane, C5H12, and hexane, C6H14. Placing methanol and ethanol, or pentane and hexane, in the bulbs shown in Figure 11.3 will result in the same diffusion and subsequent mixing of these liquids as is observed for the He and Ar gases (although at a much slower rate), yielding solutions with no significant change in energy. Unlike a mixture of gases, however, the components of these liquid-liquid solutions do, indeed, experience intermolecular attractive forces. But since the molecules of the two substances being mixed are structurally very similar, the intermolecular attractive forces between like and unlike molecules are essentially the same, and the dissolution process, therefore, does not entail any appreciable increase or decrease in energy. These examples illustrate how increased matter dispersal alone can provide the driving force required to cause the spontaneous formation of a solution. In some cases, however, the relative magnitudes of intermolecular forces of attraction between solute and solvent species may prevent dissolution.
Three types of intermolecular attractive forces are relevant to the dissolution process: solute-solute, solvent-solvent, and solute-solvent. As illustrated in Figure 11.4, the formation of a solution may be viewed as a stepwise process in which energy is consumed to overcome solute-solute and solvent-solvent attractions (endothermic processes) and released when solute-solvent attractions are established (an exothermic process referred to as solvation ). The relative magnitudes of the energy changes associated with these stepwise processes determine whether the dissolution process overall will release or absorb energy. In some cases, solutions do not form because the energy required to separate solute and solvent species is so much greater than the energy released by solvation.
Figure 11.4 This schematic representation of dissolution shows a stepwise process involving the endothermic separation of solute and solvent species (Steps 1 and 2) and exothermic solvation (Step 3).
Consider the example of an ionic compound dissolving in water. Formation of the solution requires the electrostatic forces between the cations and anions of the compound (solute–solute) be overcome completely as attractive forces are established between these ions and water molecules (solute–solvent). Hydrogen bonding between a relatively small fraction of the water molecules must also be overcome to accommodate any dissolved solute. If the solute’s electrostatic forces are significantly greater than the solvation forces, the dissolution process is significantly endothermic and the compound may not dissolve to an appreciable extent. Calcium carbonate, the major component of coral reefs, is one example of such an “insoluble” ionic compound (see Figure 11.1). On the other hand, if the solvation forces are much stronger than the compound’s electrostatic forces, the dissolution is significantly exothermic and the compound may be highly soluble. common example of this type of ionic compound is sodium hydroxide, commonly known as lye.
As noted at the beginning of this module, spontaneous solution formation is favored, but not guaranteed, by exothermic dissolution processes. While many soluble compounds do, indeed, dissolve with the release of heat, some dissolve endothermically. Ammonium nitrate (NH4NO3) is one such example and is used to make instant cold packs, like the one pictured in Figure 11.5, which are used for treating injuries. A thin-walled plastic bag of water is sealed inside a larger bag with solid NH4NO3. When the smaller bag is broken, a solution of NH4NO3 forms, absorbing heat from the surroundings (the injured area to which the pack is applied) and providing a cold compress that decreases swelling. Endothermic dissolutions such as this one require a greater energy input to separate the solute species than is recovered when the solutes are solvated, but they are spontaneous nonetheless due to the increase in disorder that accompanies formation of the solution.
Figure 11.5 An instant cold pack gets cold when certain salts, such as ammonium nitrate, dissolve in water—an endothermic process.
Link to Learning
Watch this brief video illustrating endothermic and exothermic dissolution processes.
What can speed up the dissolving process?
Describe the Dissolving Process at the Molecular Level (4.1c)
- The concept of dissolving is when a solid (called a solute) completely mixes with a liquid (called a solvent) at the molecular level to become a single visible phase and become a homogeneous mixture.
- At the molecular level, the individual solid molecules become separated and interact less with each other and more with surrounding solvent molecules.
- So a single solute molecule becomes surrounded and makes interactions with different solvent molecules.
- Random molecular motion plays into this. As the solute molecules become more separated from each other, random molecular motion will continue to separate them among the solvent molecules, thus speeding up the dissolving process.
- This is why solutes dissolve better in hot solvents versus cold solvents. Hot solvents have more kinetic energy to impart to the molecules to increase the random molecular motions.
- An example is the dissolving of sugar in water. The individual sugar molecules separate from one another and become surrounded by the water molecules. Solutes like sugar with lots of surface area dissolve better because there is more exposed surface to the solvent.
- may occur when solute particles reattach themselves after collision in a solution forming a solid
- solution is a homogeneous mixture of two or more substances.
- The component in greater quantity is called the solvent, and the component in lesser amount is called the solute
- In aqueous solutions water is the solvent.
- Solutes can be liquids, solids, or gases.
- Several terms are used to describe the degree to which a solute will dissolve in a solvent:
- When two liquids are completely soluble in each other in all proportions, they are said to be miscible.
- miscible — pairs of liquids that mix in all proportions
- For example, ethanol and water are miscible. If the liquids do not mix, they are said to be immiscible.
- immiscible — liquids that do not mix
- Oil and water, for example, are immiscible.
- A solution that contains the maximum concentration of a solute at a given temperature is called a saturated solution.
- For example, the solubility of NaCl is 6.1 moles per liter of water. An unsaturated solution has a concentration of solute that is less than the maximum concentration.
- Supersaturated solutions have a concentration of solute that is greater than that of a saturated solution
- Intermolecular forces in the solution process
- solutions tend to form when the intermolecular attractive forces between the solute and solvent molecules are about as strong as those that exist in the solute alone, or in solvent alone.
- intermolecular attraction between solute and solvent molecules is known as solvation
- Solvation is the process in which a solute particle (an ion or molecule) is surrounded by solvent molecules due to strong solute-solvent attractive forces.
- When water is the solvent the process is called hydration
- Polar and Nonpolar Solutes and Solvents
- nonpolar liquids such as heptane has intermolecular bonds with relatively weak London dispersion forces thereby making heptane immiscible in water because their attraction for each other is great due to hydrogen bond. Heptane cannot break these bonds.
- since bonds of similar strength must be broken for solvation to occur, nonpolar substances tend to soluble in nonpolar solvents, and ionic and polar substances tend to soluble in polar solvents such as water
- polar molecules are often called hydrophilic
- nonpolar molecules are called hydrophobic
- a substance that produces an electrically conducting solution when dissolved in a polar solvent, such as water
- Electrolyte solutions are normally formed when a salt is placed into a solvent such as water and the individual components dissociate due to the thermodynamic interactions between solvent and solute molecules, in a process called solvation.
- NaCl, is placed in water, the salt (a solid) dissolves into its component ions, according to the dissociation reaction
- compounds that are completely ionized in water are called strong electrolytes
- other compounds including many acids and bases, may dissolve in water without completely ionizing, thus referred to as weak electrolytes
- compounds such as glucose that dissolve with no ionization are called nonelectrolytes
Carbonation Countdown: The Effect of Temperature of Reaction Time
Have you ever wondered why bubbles form when an Alka-Seltzer tablet is dropped into water? If you’ve ever tried it, you’ve seen that the tablet fizzes furiously. The moment the tablet starts dissolving a chemical reaction occurs that releases carbon dioxide gas. This is what comprises the bubbles. Some factors can change how quickly the carbon dioxide gas is produced, which consequently affect how furiously the tablet fizzes. In this activity you’ll explore whether you can make an Alka-Seltzer tablet fizz faster or slower by changing the water’s temperature. How does this affect the reaction?
Alka-Seltzer is a medication that works as a pain reliever and an antacid. (Antacids help neutralize stomach acidity, which can cause heartburn.) The pain reliever used is aspirin and the antacid used is baking soda, or sodium bicarbonate. The tablets also include other ingredients, such as citric acid (a weak acid that adds flavor—as well as provides important hydrogen ions, which will come into play as you shall soon see).
To take the tablets, they’re fully dissolved in water, where they famously undergo a chemical reaction that produces lots of carbon dioxide bubbles—or fizz. Why is this? As the tablets dissolve, the sodium bicarbonate splits apart to form sodium and bicarbonate ions. The bicarbonate ions react with hydrogen ions from the citric acid to form carbon dioxide gas (and water). This is how the bubbles are made.
How is temperature related to this reaction? For the reaction to occur, the bicarbonate ions must come into contact with the hydrogen ions in just the right way. The probability of the bicarbonate and hydrogen ions doing this is affected by temperature: the higher the temperature, the faster the molecules move; the lower the temperature, the slower they move. (The temperature of a solution is a measure of its molecules’ average motion and energy.) Can you guess whether fast-moving molecules or slow-moving ones will speed the reaction time?
• Two identical jars (You can also use drinking glasses, clear plastic cups, bottles or vases.)
• Enough ice cubes to fill one of the jars halfway
• Cold tap water
• Hot tap water
• Two Alka-Seltzer tablets
• Timer or clock that shows seconds
• Optional: helper
• Fill one of the jars halfway with ice cubes. Add cold tap water to about an inch from the rim. Stir the ice cubes in the jar for about a minute so that the temperature evens out. Right before you start the activity use a spoon to remove the cubes.
• Add hot tap water to the second, empty jar until it is about an inch from the rim. Be careful when handling the hot water.
• Continue with the procedure immediately after preparing the jars (so that the water in the jars is still very cold or very hot).
• Drop an Alka-Seltzer tablet into the jar with hot water. Time how long it takes for the tablet to disappear. You may want to have a helper time the reaction. How long does it take the tablet to disappear? How vigorous are the bubbles?
• Drop an Alka-Seltzer tablet into the jar with the ice-cold water (after having removed the ice cubes with a spoon). Again time how long it takes the tablet to disappear. How long does it take the tablet to disappear in the colder water?
• Do you notice other differences in how the reaction happens in the colder versus in the hotter water?
• Why do you think you got the results you did?
• Extra: Test Alka-Seltzer tablets in a wider range of temperatures, and then draw a graph showing the time it takes a tablet to dissolve in water at each temperature (check with a thermometer). What temperature change is required to increase the reaction time by a factor of two (make it as twice as fast)? What about decreasing the reaction time by a factor of two?
• Extra: Compare whole Alka-Seltzer tablets to pieces of Alka-Seltzer tablets. If there is a greater surface area (that is, a tablet is broken up into more pieces to expose more surface), does the same amount of tablet result in the reaction happening faster or slower?
• Extra: You can turn this activity into a homemade lava lamp! To do this, you will use an empty container, such as a tall jar or clear plastic one- or two-liter bottle. Fill it with about two inches of water, add five drops of food coloring and then fill it at least three quarters full with vegetable oil before adding one quarter of an Alka-Seltzer tablet. You could repeat this activity using your homemade lava lamp at colder and warmer temperatures. (Because it contains oil, you should have an adult help you devise a safe way to warm or cool the contents of each container.) How does the bicarbonate reaction look in the homemade lava lamp?
Observations and results
Did the Alka-Seltzer tablet dissolve much faster in the hot water compared to the cold? Were there a lot more bubbles produced initially in the hot compared with the cold water?
After the Alka-Seltzer tablet was added to the hot water the tablet should have quickly dissolved, taking some 20 to 30 seconds to do so, depending on the exact temperature. After the tablet was added to the ice-cold water it should have taken much longer to dissolve, with most of the tablet disappearing after about two to three minutes, but with some bubbles still apparent after six minutes or longer. In the hot water the tablet should have more vigorously produced bubbles than in the cold water. The higher the temperature, the faster the molecules move—and the more likely it is that the bicarbonate will contact hydrogen in just the right way for the chemical reaction to occur and produce carbon dioxide bubbles.
This activity brought to you in partnership with Science Buddies