Solutions for the DAT
/Learn key DAT concepts about solutions, plus practice questions and answers
everything you need to know about solutions for the dat
Table of Contents
Part 1: Introduction to solutions
Part 2: Solubility rules
a) Solubility rules
b) Concentration units
c) Precipitation reactions
d) Common ion effect
Part 3: Colligative properties
a) Boiling point elevation and freezing point depression
b) Vapor pressure depression
c) Osmosis and osmotic pressure
Part 4: High-yield terms
Part 5: Questions and answers
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Part 1: Introduction to solutions
Solutions, a fundamental concept for the DAT, are homogeneous mixtures where one or more substances are uniformly distributed in another substance. The study of solutions encompasses a range of crucial topics, including concentration, colligative properties, and factors influencing solubility. Understanding the behavior of solutions is essential in various scientific and everyday contexts, such as chemical reactions, biological processes, and environmental phenomena. As you study this guide, pay attention to the bolded words. These are high-yield terms that you need to know for the DAT.
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Part 2: Solubility rules
a) Solubility rules
Solubility pertains to a solute's capacity to dissolve in a specific solvent. A solute is a substance that is dissolved into another substance, known as the solvent. When a solute is in equilibrium between its dissolved and undissolved states, the solution is deemed saturated. Any additional solute introduced becomes insoluble, leading to the formation of a precipitate, representing the solute's solid, undissolved state—a phenomenon known as precipitation reactions.
The solubility of a substance is influenced by the solvent used and the temperature of the solution. Polar solvents, like water, have an affinity for dissolving polar or ionic compounds. Nonpolar solvents, such as oil or fat, favor the dissolution of nonpolar compounds, like vitamin A.
In the context of the DAT, water is the most commonly encountered solvent. The following table summarizes key solubility rules crucial for test day, though it's important to note that exceptions exist for each of these rules.
| Water Solubility Rules |
|---|
| Salts containing Group I elements (alkali metals) are soluble |
| Salts with the ammonium ion (NH4+) are soluble |
| Salts with the acetate ion (CH3COO-) are soluble |
| Salts with the nitrate ion (NO3-) are soluble |
|
Salts containing the sulfate ion (SO42-) are soluble
Exceptions: SrSO4, PbSO4, BaSO4, Ag2SO4, CaSO4 |
|
Sulfites (SO32-) are insoluble
Exceptions: Sulfites with Group I elements and ammonium |
|
Sulfides (S2-) are insoluble Exceptions: Sulfides with Group I elements and ammonium |
|
Halides are soluble (e.g., Iodine, Chlorine, Bromine) Exceptions: Fluoride and halides containing Pb2+, Ag+, and (Hg2)2+ |
|
Hydroxide salts with ammonium, Group I elements (alkali metals), and certain Group II elements (Ca2+, Sr2+, and Ba2+) are soluble All other hydroxide salts are insoluble |
|
Phosphates (PO43-) are insoluble Exceptions: Phosphates containing Group I elements and ammonium |
|
Carbonates (CO32-) are insoluble (e.g., CaCO3, FeCO3, and SrCO3) Exceptions: Carbonates containing Group I elements and ammonium |
Memorizing this compilation of solubility rules might seem daunting, but there are general guidelines. Compounds like ammonium, acetate, nitrate, and sulfate are typically soluble. On the other hand, sulfites, sulfides, calcium, and compounds involving transition metals are usually insoluble in water.
b) Concentration units
In the process of creating a solution, a solute is dissolved in a solvent, typically involving a solid compound placed into a liquid-phase solvent. The amount of solute within a specific volume of solvent is termed concentration. Solutions with a lower ratio of solute to solvent are classified as dilute, while those with a higher ratio are deemed concentrated. Various units can express the concentration of a solution, with molarity (M) being the most common. Molarity represents the moles of solute per liter of solution. Molality (m) is an alternative unit, indicating the moles of solute per kilogram of solvent. It is essential to note that molarity and molality, while sharing similar symbols, have distinct meanings.
Osmoles, referring to the moles of individual solute particles, are measured in terms of osmolarity (osmoles per liter of solution) and osmolality (osmoles per kilogram of solvent). Another unit is normality (N), denoting equivalents per liter of solution, where equivalents can represent specific species of interest. For example, in the dissociation of an acid, the relevant quantity is often the concentration of hydrogen ions. Mole fraction offers an alternative expression for concentration, representing the ratio of moles of one substance to the total moles within a solution or mixture.
| Unit | Symbol | Formula |
|---|---|---|
c) Precipitation reactions
In the dissolution of compounds within a solution, replacement or synthesis reactions can give rise to new compounds. If such a reaction yields an insoluble compound, the resultant product will take on a solid form and precipitate out of the reaction. These types of precipitation reactions are shown below.
FIGURE 1: COMMON TYPES OF PRECIPITATION REACTIONS
In chemical reactions, the net ionic equation represents the essential changes occurring in a reaction by omitting spectator ions. Spectator ions are ions that do not participate in the chemical change and thus appear unchanged on both sides of the equation. The net ionic equation focuses solely on the ions involved in the actual chemical transformation, providing a clearer picture of the fundamental processes taking place. For instance, consider the reaction between silver nitrate (AgNO3) and sodium chloride (NaCl) to form silver chloride (AgCl) precipitate. The complete ionic equation for this reaction would include all ions present before and after the reaction, such as Ag+, NO3−, Na+, and Cl−. However, by recognizing that Na+ and NO3+ ions are spectator ions that do not participate in forming the precipitate, the net ionic equation focuses solely on the key change:
Ag+(aq)+Cl−(aq)→AgCl(s)
Conversely, a complete ionic equation includes all ions, both reactant and product, in their ionic forms. Unlike the net ionic equation, it does not differentiate between ions that undergo chemical change and spectator ions. In the example above, the complete ionic equation would include all ions present in the solution before and after the reaction, illustrating the dissociation of all soluble compounds into their constituent ions. Therefore, while the complete ionic equation provides a comprehensive view of all ions involved in a reaction, the net ionic equation focuses specifically on the ions directly involved in the chemical transformation, simplifying the representation of the reaction's essential components.
d) Common ion effect
Apart from the previously mentioned factors, solubility is further influenced by the common ion effect. This phenomenon refers to the decrease in the solubility of a salt when introduced into a solution containing one of its ions. To comprehend this effect, it is crucial to grasp the dissociation reaction. Considering a generic solute represented by the chemical formula AaBb, the dissociation reaction can be expressed as follows:
AaBb (s) ⇋ aA (aq) + bB (aq)
Here, A and B denote the constituent ions. When a certain quantity of these ions is already present in the solution, le Chatelier’s principle comes into play upon the addition of extra reactant. The existence of the dissociated ion or product prompts an equilibrium shift to the left. Consequently, any additional solute is more likely to remain in its undissociated form.
For instance, if potassium chloride (KCl) is dissolved in a solution containing potassium ions, its solubility will be lower compared to when it is added to pure water. The common ion effect proves advantageous in laboratory separations, where the formation of complex ions, often metallic or insoluble, leads to the precipitation of a solid. This precipitate can then be easily separated from the aqueous solution and subsequently dried.
