Solutions and Aqueous Ionic Reactions

General Terminology

A solution is a mixture that is homogeneous on the particle level, meaning that the solute particles are evenly distributed among the solvent particles. The solute is the substance that is dissolved (present in the smaller amount), and the solvent is the substance that does the dissolving (present in the greater amount). A substance is described as soluble if it will dissolve in a particular solvent. For instance, water soluble means the substance will dissolve in water. When the solvent and solute (s) are both (or all) liquids, they can typically be mixed in any proportion (such as with water and alcohol). In such cases, the substances are described as miscible. Liquids that will not mix, such as water and oil, are often described as immiscible. The amount of solute that will dissolve in a given amount of solvent is called its solubility (e.g. at a certain temperature, the solubility of sucrose in water is 240.0 g/100 g H₂O.

The dissolving process and heat of solution - 4.1 & 11.2

The process is called solvation. For a solid dissolving in a liquid, solvent particles are attracted to solute particles in the crystal. They separate from each other (an endothermic process) and begin surrounding the solute particles (an exothermic process), pulling them away from the other solute particles (another endothermic process). Individual solute particles are surrounded by several solvent particles (solvated). (When water is the solvent, the terms hydration and hydrated can be used in place of solvation and solvated.)

If the heat required to separate the solute particles from each other plus the heat required to separate the solvent particles from each other is greater than the heat released when the two types of particles mix, then the overall heat of solution (∆H_soln - the heat associated with the dissolving process) will be endothermic. If the heat required to separate the solute particles from each other plus the heat required to separate the solvent particles from each other is less than the heat released when the two types of particles mix, then the overall heat of solution will be exothermic

Electrolytes and Nonelectrolytes - 4.2

Water is a nonconductor of electricity. Some solutes dissolve to produce solution that will conduct electricity. Such solutes are called electrolytes. In an electrolyte, some or all of the solute particles break apart to form ions. The presence of these ions (charged particles) allows for electricity to be conducted through the solution. The process of forming ions while dissolving is called dissociation. The process of dissociation is represented by an equation for dissociation. The equation shows how the compound breaks apart into ions. For example

NaCl → Na¹⁺ + Cl^1-

K₂SO₄ → 2 K¹⁺ + SO₄^2-

Fe₃(PO₄)₂ → 3 Fe²⁺ + 2 PO₄^3-

Strong electrolytes dissociate completely as they dissolve. For weak electrolytes, only a small percentage of the dissolved substance experiences dissociation. There a three types of substances that behave as electrolytes: acids (like HCl), bases( like NaOH), and salts (like NaCl).

Nonelectrolytes dissolve but do not dissociate. All of the solute particles are neutral, so electricity cannot flow through the solution. Sugar is the most familiar nonelectrolyte.

Concentration: Units (molarity), preparing solutions, measuring solutions - 4.3

There are general qualitative terms used to describe the amount of solute present: dilute (a little solute), concentrated (a lot of solute), and saturated (the maximum amount of solute). Note that a saturated solution may not be concentrated because some solute are only slightly soluble.

For quantitative applications-when the precise amount of solute needs to be known-concentration units are required. The most common concentration unit is molarity. Molarity is the number of moles of solute per liter of solution: M = n/V.

To prepare a solution of a particular molarity, the solute is measured and placed in a volumetric flask. The flask is then filled half-way with solvent. After the solid has dissolved, additional pure solvent is added to fill the volumetric flask (to the mark on its neck).

When preparing a solution of a solid solute in water, the amount of solid required can be determined by multiplying the desired molarity times the desired volume to find the number of moles of solute required (MV = n). The number of moles is then multiplied by the molar mass of the solute to find the mass in grams. (n×MM = m).

When preparing a solution from a more concentrated solution, the volume of concentrated solution required can be determined using the equation M_cV_c = M_dV_d (the ‘c’ subscript indicates the concentrated solution, the ‘d’ indicates the dilute solution).

Ionic Reactions - 4.4

Ionic reactions are those in which ions do not change charge but simply change partners. Double displacement reactions are all ionic. Two specific types are precipitation reactions, in which a solid product is formed by the reaction of two aqueous reactants and acid-base neutralization reactions, in which an acid reacts with a base to produce a salt and water.

Precipitation reactions, solubility rules - 4.5

When two aqueous reactants are mixed and one of the products is insoluble in water, that compound forms a solid. This solid is called a precipitate. To determine whether an ionic compound is soluble, the solubility rules are used. The insoluble product will be the precipitate.

Ionic and net ionic equations - 4.6

For ionic reactions, ionic and net ionic equations can be written. The ionic equation shows strong electrolytes (soluble ionic compounds and strong acids) in their dissociated form while leaving the formulas of weak acids, solids, liquids, and gases unchanged. The net ionic equation removes the spectator ions (the ions that appear in exactly the same form on both sides of the equation), leaving only the species that are actually changed during the reaction.

Formula equation: 2 Na₃PO₄ (aq) + 3 CaCl₂ (aq) → Ca₃PO₄)₂ (s) + 6 NaCl (aq)

Ionic equation: 6 Na⁺ + 2 PO₄^3- + 3 Ca²⁺ + 6 Cl^1- → Ca₃(PO₄)₂ (s) + 6 Na⁺ + 6 Cl^1-

Net ionic equation: 2 PO₄^3- + 3 Ca²⁺ → Ca₃(PO₄)₂ (s)

(The spectator ions were Na⁺ and Cl^1-.)

Stoichiometry in solutions - 4.7

When working with solutions, it is more convenient to measure volumes than masses. Consequently, it is more convenient to use molarities for stoichiometric calculations than it is to used masses. The basic process of stoichiometry remains the same:

1. Convert the given information to moles

2. Find the common factor.

3. Find the moles of the desired substance.

4. Convert moles to the desired form.

Now there are two ways to find mols:

If given the mass of a compound, divide by its molar mass.

If given volume and molarity ofr a solution, multiply them ((be sure volume is in liters).

At the end of the process, there are three possible ways to express the amount of the desired substance:

If asked to find mass, multiply the number of moles of the desired substance by its molar mass.

If asked to find the volume of a solution, divide the number of moles by the molarity (which will be given in the problem).

If asked to find the molarity of a solution, divide the number of moles by the volume (which will be either be given in the problem or can be found (for a product) by adding the reactant volumes).

Gravimetric analysis - not in text

Gravimetric analysis is a type of quantitative analysis (any procedure used to determine the specific amount of a substance that is present). Excess solution is added to the solution being tested in order to precipitate all of the desired ion. The solid is then filtered from the solution, dried, and weighed. The mass collected can be used to calculate the concentration in the original solution. For example, in a previous lab we added excess sodium carbonate to precipitate all the Ni²⁺ or Cu²⁺ in a solution of unknown concentration.

Acid-base reactions, titration - 4.8

An acid is substance that produces hydrogen ion (H⁺) in water. Aqueous hydrogen ion can also be represented as the hydronium ion (H₃O¹⁺). Some acids are monoprotic (like HCl), meaning each acid molecule produces one hydrogen ion in solution (remember, a hydrogen ion is a proton). Acids can also be polyprotic (like H₂SO₄ and H₃PO₄), meaning each acid molecule produces more than one hydrogen ion in solution (poly means ‘many’). H₂SO₄ is diprotic (2 H⁺’s per molecule), and H₃PO₄)is triprotic (3H⁺’s).

Acids consist of covalently bonded molecules, but they are highly polar and will ionize in water. This means all acids are electrolytes; a few are strong (dissociating completely), but most are weak (dissolving readily but dissociating only a little). There are only a few strong acids: HCl, HBr, HI, HNO₃, HClO₄, and H₂SO₄ (first dissociation only).

A base is a substance that produces hydroxide ion (OH^1-) in water. While there are other types of bases, we will consider just the metal hydroxides (like NaOH and Ca(OH)₂). The term alkaline used synonymously with the term basic.

An acid and a base will react to produce a salt and water. A salt is an ionic compound formed by the reaction of an acid with a base. The process is called neutralization because the product is much closer to being neutral than the original acid or base.

Some substances, such as water, are amphoteric, meaning they can behave as either an acid or a base depending on the circumstances.

Titration is a lab technique for determining the concentration of a solution by reacting it with a solution of known concentration. Acid-base titration is one of the more common types. A basic solution of known concentration is placed in a burette. The substance in the burette is called the titrant. An acidic solution of unknown concentration is placed in an Erlenmeyer flask. The volume must be precisely measured, and an indicator is added to the solution. An indicator is a substance that changes color depending on the pH.

The base is then added dropwise from the burette until the acid solution permanently changes color. One drop of titrant should cause the color change. This point in the experiment, the point at which you stop the titration, is called the endpoint. The total volume of titrant used can be read directly from the burette. Because you know both the volume and the molarity of the base as well as the volume of the acid, the balanced chemical equation can be used to determine the molarity of the acid. You must assume that the endpoint is very close to the equivalence point (the theoretical point at which the acid and base exactly neutralize each other. (A titration could also be done using an acidic solution of known concentration as the titrant to find the molarity of a basic solution).

Ionic and net ionic equation can also be written for acid base reactions. It is important to remember that only the strong acids dissociate significantly, so weak acids are not separated into ions in the ionic forms of the equation. Only aqueous salts and strong acids are represented as dissociated. An example of a reaction involving a weak acid:

_{(weak acid)}

Formula equation: 2 NaOH (aq) + H₂SO₃ (aq) → Na₂SO₃ (aq) + 2 H₂O (l)

Ionic equation: 2 Na⁺ + 2 OH^1- + H₂SO₃ (aq) → 2 Na⁺ + SO₃^2- + 2 H₂O (l)

Net ionic equation: 2 OH^1- + H₂SO₃ (aq) → SO₃^2- + 2 H₂O (l)

An example of a reaction involving a strong acid:

Formula equation: LiOH (aq) + HNO₃ (aq) → LiNO₃ (aq) + H₂O (l)

Ionic equation: Li¹⁺ + OH^1- + H¹⁺ + NO₃^1- (aq) → Li¹⁺ + NO₃^1- + 2 H₂O (l)

Net ionic equation: OH^1- + H¹⁺ → H₂O (l)

Occasionally the base or acid may be in the solid state in the reaction and so would not be shown as dissociated.

Oxidation-Reduction Reactions (Chapter 4)

Know how to:

Assign oxidation numbers to free elements, elements in a compound, and elements in a polyatomic ion

Determine whether a chemical reaction is a redox reaction or not

Identify oxidizing agents and reducing agents

Balance a redox reaction in acidic or in basic solution using the half cell method

Redox reactions, or oxidation-reduction reactions, are reactions in which electrons are moved from one atom to another.

Oxidation involves the loss of electrons, resulting in an increase in the oxidation state (Zn ® Zn²⁺ + 2 e^-).

Reduction involves gaining electrons, causing a decrease in the oxidation state (I₂ + 2 e^- ® 2 I^-).

Oxidation and reduction must always occur together so electrons can be conserved. Single displacement, combustion, and synthesis reactions are all redox reactions. In addition, most decomposition reactions and many other types of reactions are also redox.

The reactant that causes oxidation is the oxidizing agent. The oxidizing agent contains the element that experiences reduction during the reaction.

The reactant that causes reduction is the reducing agent. The reducing agent contains the element that experiences oxidation during the reaction.

In the reaction 4 Al + 3 O₂ ® 2 Al₂O₃ , the oxidation state of Al increases from zero to +3, so it is experiencing oxidation, and the Al metal is the reducing agent. The oxidation state of oxygen changes from zero to -2, so it is experiencing reduction, and the O₂ gas is the oxidizing agent. The reaction does not need to be balanced to identify these agents.

A redox reaction can be divided into two half-reactions. For the reaction above, the oxidation half reaction would be Al ® Al³⁺ + 3 e^-, and the reduction half reaction would be O₂ + 4 e^- ® 2 O^2-. To conserve electrons, the oxidation half reaction would need to happen 4 times as the reduction is happening 3 times (electrons lost must equal electrons gained, in this case 12 in each half).