Solutions

Overview: In this module you will learn the vocabulary that chemists use in describing solutions as well as how to calculate solution concentrations.

Skills:

• Calculation of molarity, molality and mole fraction
• Interconversion between molarity, molality and mole fraction
New terms:
 Mixture Compound Solution Solvent Solute Aqueous Solution Solid Solution Molarity Molality Mole Fraction

A mixture is matter that contains more than one pure substance and can be separated into its components by using physical techniques. The composition is variable and its properties are related to its individual components. In a mixture, the components retain their own chemical properties. An example of a mixture is the combination of rice and salt. Compare this to a compound.

A compound is a substance that contains two or more different elements with their atoms in a definite ratio. Compounds cannot be separated by physical techniques such as filtering. The composition is the same throughout. Water will always have two hydrogen atoms and one oxygen atom. If instead the ratio were two hydrogen atoms to two oxygen atoms, then the compound is no longer water (H2O), it is now hydrogen peroxide (H2O2). Elements in a compound are not just mixed together; they are bonded together in a specific way. The properties of a compound are usually very distinct from the properties of the individual elements that make the compound. For instance, sulfur, a yellow solid, combines with oxygen, an oderless gas, to form sulfur dioxide (SO2), which is a poisonous, colorless, pungent gas.

A solution is a special type of mixture that is homogeneous throughout. This means that the molecules or ions involved are so well mixed that the composition is uniform throughout the mixture. Think of salt-water. You cannot see salt within the water when it is fully dissolved, not even with the aid of a microscope. (This contrasts with a heterogeneous mixture in which you can identify the separate components. For example, a mixture of salt and sand is heterogeneous.)

A solvent is the component in a solution that is present in the largest amount. In a NaCl solution (salt-water), the solvent is water.
A solute is the component in a solution in the lesser amount. In a NaCl solution, the salt is the solute. A solution may contain more than one solute.

There are different types of solutions. The one you are probably most familiar with is the aqueous solution.
An aqueous solution is a solution in which water is the solvent. A NaCl solution is an aqueous solution.
A non-aqueous solution is a solution in which water is not the solvent. Examples of non-aqueous solutions are solutions used in dry cleaning (a solution of ethene in the solvent dichloromethane).
A solid solution is a solution in which a solid is the solvent. An example is a brass solution that is formed by dissolving copper in zinc.

So what happens when you drop salt into a glass of water? The water before and after does not look different (assuming that all of the salt is dissolved). However, if you took a drink of it, it certainly tastes different. It boils at a higher temperature than pure water and it conducts electricity. What happened?

Your everyday table salt consists of NaCl. Water is made of H2O molecules. When these two are combined together in a solution, NaCl actually separates into ions. [In solid NaCl, Na+ and Cl- ions are arranged in an ordered three dimensional array called a crystal lattice, as depicted in the figure below.] Thus NaCl (s) becomes Na+ and Cl- ions in solution (i.e., NaCl dissolves in water). Why does NaCl dissolve?

NaCl becomes solvated ions because of favorable electrostatic interactions (you will learn about this in Chemistry 111) and favorable entropy conditions (you will learn about this in Chemistry 112). In order for the solute to dissolve, these two effects must be stronger than the interactions within the NaCl crystal and the solvent molecules with themselves. In other words, NaCl dissolves in water because the electrostatic interactions and the entropic effects are stronger for the ions and water than they are for the NaCl crystal and the H2O by itself. It is okay if you do not fully understand this concept yet, it will become clearer throughout general chemistry.

Even though NaCl dissolves to become ions in a solvent, the overall charge remains neutral. Remember that it is NaCl, a neutral compound, that forms the Na+ and Cl- ions. There will be an equal number of positive and negative charges; therefore, the solution will be neutral.

 Water NaCl Crystal Lattice  Water is a polar molecule. This means that the charges are not evenly distributed. The oxygen atom has a partial negative charge (d-) and the hydrogen atoms have a partial positive charge (d+).

The following reaction represents the dissolution of NaCl (s) in water:  The diagrams above show the dissolution of NaCl solid. The water molecules surround the ions. These ions are now free to move about in solution since they are no longer in a crystal lattice. This means that the ionic bonds between Na+ and Cl- are broken.

Notice that the oxygen of H2O surrounds the Na+ and the hydrogens of H2O surround the Cl-. This has to do with the polar properties of water and the charges on the ions. The partial negative charge on oxygen is attracted to the positive charge on the Na+ ion (opposites attract). The same is true for the hydrogen; the partial positive charge on hydrogen is attracted to the negative Cl- ion. You will learn about polarity and intermolecular interactions in Chemistry 111.

Measuring a Solution

When discussing solutions, we typically talk about the solution's concentration. In chemistry, we use molarity to calculate the concentration. Other important terms are the molality and mole fraction of a solution.

The molarity is the number of moles of solute per liter of solution. This is a specific concentration measurement. Molarity is defined as the number of moles of solute per unit volume. The molarity is reported as M (read molar), which is mol of solute/L of solution. Molarity is temperature dependent as the volume of the density of a solution typically changes with temperature.

The molality is the number of moles of solute per kilogram of solvent. This measurement is not temperature dependent, as the mass does not change with temperature. The units are denoted by m, which is read as molal and is mol of solute/kg of solvent.

A mole fraction, as the name implies, is a comparison of the number of moles in solution. It is found by taking the number of moles of solutes (or solvent) divided by the total number of moles (solutes + solvent). Since this is a fraction, there are no units. The mole fractions of a solution must add up to one.

Let's look at a simple solution made of two components, 1 and 2. X is the mole fraction and n is the number of moles. Example 1.

A solution is prepared by dissolving 34.2 g of MgCl2 in 0.430 L of H2O. Calculate the molarity, molality and mole fraction of MgCl2 if the density of water is 1.00 g/cm3 and the density of the solution is 1.089 g/cm3.

• Mole Fraction: For MgCl2 we need to calculate the number of moles by using the molecular mass, 95.211 g/mol. For water, the volume was given as well as the density. We can use these two values to determine the mass and use the molecular mass to determine the number of moles of H2O. The problem asks for the mole fraction of MgCl2. • Molarity: We must know the number of moles of solute, which we just calculated, and the volume of the solution in liters. To determine the volume of solution, use the given density. In order to use the density, however, we first need to find the total mass of the solution. • Molality: To find the molality, we need the number of moles of solute and the mass (in kg) of solvent, both of which have been previously calculated. Notice in the example that the molarity and molality are very close to each other. This will be true in an aqueous solution as the density of water is very nearly 1 g/mL. This relationship does collapse in a non-aqueous solution and in highly concentrated aqueous solutions.

Example 2.

For the solution in Example 1, 34.2 g of MgCl2 in 0.430 L of H2O (r = 1.089 g/cm3), calculate the molarity, molality and mole fraction of the Cl- ion in solution.

Now remember what is happening. When the MgCl2 is placed in water, it dissolves into Mg2+ and Cl- ions.

• Mole Fraction: Again we will start with the mole fraction. To determine this, the number of moles of Cl- is needed. We already calculated the number of moles of water in Example 1, 23.9 mol. However, the mole fraction is the number of moles for Cl- divided by the total number of moles of all species in the solution. So we also need to know the number of moles of Mg2+ in solution. Now the mole fraction is calculated by: • Molarity: For the molarity, the number of moles of solute per liter of solution is needed. Fortunately we have already calculated these quantities. • Molality: Finally for the molality, we need the moles of solute and mass of solvent. We also have calculated these numbers before. Notice that the molarity and molality are twice as much for the Cl- ion as for MgCl2. This is because 1 mole of MgCl2 dissociates producing 2 moles of Cl- in the aqueous solution.

Molecular Solutions

In molecular solutions, bonds are not broken as they are in non-molecular solutions (also sometimes called ionic solutions). NaCl (aq) is an example of a non-molecular solution. Recall that in non-molecular solutions the ionic bonds were broken within the compound. For molecular solutions, Glucose, a sugar molecule, is an example of a compound that forms a molecular solution in water. In the solid, crystalline form glucose molecules are also ordered into a three dimensional array, as in the case of the NaCl crystal lattice discussed above. However, unlike in the case of NaCl, in which the compound breaks apart into smaller components (the individual ions), the glucose molecule remains intact as a single molecular unit in solution. The figure below is a schematic of the glucose molecule dissolved in water. Each glucose molecule is surrounded by a certain number of water molecules; therefore in solution, the glucose solid broke apart, but the glucose molecules themselves remained intact. Note that the definitions of molarity, molality, and mole fraction are the same for both molecular and nonmolecular solutions. You will learn more about interactions between molecules in Chemistry 111. A schematic of the glucose molecule in aqueous solution.

Advanced Applications: Washington University chemists synthesize miniature electronic components for nanotechnology.

Summary

Now you should have a good understanding of solutions and the solvation process for ions in aqueous solutions. You need to be comfortable with using molarity, molality and mole fractions. Also you should know the definitions pertaining to solutions.

Practice Quizzes: Stoichiometry/Limiting Reagents/Solutions These two quizzes cover the three tutorial modules Stoichiometry, Limiting Reagents, and Solutions. You will probably want to review all three of these modules before trying the quiz.

Note: You will need a pencil, scratch paper, calculator, periodic table and equation sheet to work the practice quiz. Quiz questions are timed (4 minutes per question).

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Open Equations/Constants Page in a separate browser window.