Aqueous Solutions and Net Ionic Equations

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3.3 Aqueous Solutions and Net Ionic Equations

For a chemical reaction to occur, the reactants involved must be able to come into contact with one another. In the gas or liquid phase, the molecules that make up the reactants move readily, allowing this contact to occur. In the solid phase, though, such motion is uncommon and reactions occur only very slowly if at all. One way to en- able the needed contact between reactants that are solids under normal conditions is to dissolve them. Water is the most common medium for producing such solutions.

Reactions that occur in water are said to take place in aqueous solution. To describe this important group of chemical reactions, fi rst we need to defi ne some key terms.

Solutions, Solvents, and Solutes

Although water accounts for the vast majority of all the molecules in an aqueous solution, it is neither a reactant nor a product in many aqueous reactions. Water simply serves as the medium in which the reaction occurs. The entire liquid is called a solution, meaning

CH₂O + O₂ CO + H₂O

Change CO to CO₂?

Correct balanced form Correct species, balanced equation, valid molecular interpretation

Undesirable Correct species, balanced equation, no valid molecular interpretation

Not valid Correct species unbalanced equation Not valid

Incorrect species, unbalanced equation

Correct species,

unbalanced equation Multiply O₂ by ¹/₂?

Multiply O₂ by 2?

Multiply ALL coefficients by 2?

CH₂O + O₂ CO₂ + H₂O

CH₂O + O₂ CO + H₂O

2 CH₂O + O₂ 2 CO + 2 H₂O

CH₂O + ¹/₂O₂

?

CO + H₂O

Figure 3.5 ❚ The molecular drawings here illustrate some common errors in balancing chemical equations.

that it is a homogeneous mixture of two or more substances. In any solution, the compo- nent present in greater amounts—usually much greater amounts—is called the solvent.

In an aqueous solution, water is the solvent. Minor components of the solution are called solutes. The key feature of a solution is that the solutes dissolve in the solvent. When a substance dissolves, its particles are dispersed in the solvent, and typically many solvent molecules will surround any individual molecule or ion of solute. Although water is the most common solvent, it is not the only one. Not all solutions are even liquids. The same ideas can be applied to homogeneous mixtures of gases or solids. Air is a good example of a gaseous solution, and alloys like brass can be described as solid solutions.

If we want to describe a pure substance, all we need to specify is its identity. If you pick up a bottle in the chemistry lab that is labeled “distilled water,” for example, you would know exactly what is in the container. But if the same container says only that it contains a solution of salt (NaCl) dissolved in water, you can’t be as sure what the contents are. Specifi cally, you would probably want to know something about the relative amounts of salt and water that were mixed to prepare the solution. So we need to specify at least one other key piece of information: the concentration. If there are many solute particles present, the solution is said to be concentrated. When there are few solute particles present the solution is called dilute. In Section 3.5, we will discuss various units that can be used to report solution concentrations.

Some solutes can be dissolved in water to produce very concentrated solutions, whereas other substances do not dissolve to any measurable extent. So we can charac- terize various compounds in terms of their solubility in water. Those compounds that dissolve readily are said to be soluble, whereas those that do not dissolve signifi cantly are called insoluble. Chemists often use a series of solubility rules to predict whether or not a particular compound is likely to dissolve in water. The most common rules are summarized in Table 3.1.

In Chapter 12, we will see that solubility is not an “either or” process, but that there is continuous variability. These rules summarize the main quantitative tendencies we will learn then.

Table

❚

^3.1

Solubility guidelines for ionic compounds in water at room temperature

Usually Soluble Exceptions

Group 1 cations (Li⁺, Na⁺, K⁺, Rb⁺, Cs⁺), ammonium (NH4+)

No common exceptions

Nitrates (NO3−), nitrites (NO2−) Moderately soluble: AgNO2

Chlorides, bromides, iodides (Cl⁻, Br⁻, I⁻)

Insoluble: AgCl, Hg2Cl2, PbCl2, AgBr, Hg2Br2, PbBr2, AgI, Hg2I2, and PbI2

Fluorides (F⁻) Insoluble: MgF₂, CaF₂, SrF₂, BaF₂, PbF₂ Sulfates (SO42−) Insoluble: BaSO4, PbSO4, HgSO4

Moderately soluble: CaSO4, SrSO4, Ag2SO4

Chlorates (ClO3−), perchlorates (ClO4−) No common exceptions

Acetates (CH3COO⁻) Moderately soluble: AgCH3COO

Usually Insoluble Exceptions

Phosphates (PO43−) Soluble: (NH4)3PO4, Na3PO4, K3PO4

Carbonates (CO32−) Soluble: (NH4)2CO3, Na2CO3, K2CO3

Hydroxides (OH⁻) Soluble: LiOH, NaOH, KOH, Ba(OH)2

Moderately soluble: Ca(OH)2, Sr(OH)2

Sulfi des (S²⁻) Soluble: (NH4)2S, Na2S, K2S, MgS, CaS Figure 3.6 ❚ In this sequence of

photos, one of the authors prepares aqueous solutions of CuSO4. In the upper left panel, solid CuSO4!the solute!is transferred to a fl ask. In the upper right panel, water!the solvent!is added. The fl ask is shaken to speed the dissolution process (lower left). The fi nal photo shows two CuSO4 solutions of different concentrations. The solution on the left has the higher concentration, as seen from its darker color.

Photos: Lawrence S. Brown

E X A M P L E P RO B L E M 3 . 2

Which of the following compounds would you predict are soluble in water at room temperature? (a) KClO₃, (b) CaCO₃, (c) BaSO₄, (d) KMnO₄

Strategy Solubility guidelines for common ions are given in Table 3.1. So we will identify the ions in each compound and consult the table as needed to determine the solubilities.

Solution

(a) KClO₃ is potassium chlorate. From the solubility guidelines in Table 3.1, we see that compounds containing K⁺ and ClO₃⁻ tend to be soluble and that no common exceptions are mentioned. So we predict that KClO₃ should be soluble.

(b) CaCO₃ is calcium carbonate. Again consulting the table, we see that carbonates are generally insoluble and that CaCO₃ is not listed among the exceptions. Thus CaCO₃ should be insoluble.

(c) BaSO₄ is barium sulfate. Although most sulfates are soluble, BaSO₄ is listed in the table as an exception to that rule. Therefore we expect BaSO₄ to be insoluble.

(d) KMnO₄ is potassium permanganate. Although the permanganate ion (MnO₄⁻) is not listed in Table 3.1, the table does tell us that all compounds of K⁺ are soluble.

So we would predict that KMnO₄ should be soluble.

Discussion Here we simply consulted the table to check the solubility of each compound. Chemists generally gain familiarity with many of these solubility rules and learn to recognize soluble and insoluble salts without consulting such a table. You should check with your instructor to see whether you are expected to memorize these rules and exceptions.

Check Your Understanding Which of the following compounds would be soluble in water at room temperature? (a) NH₄Cl, (b) KOH, (c) Ca(CH₃COO)₂, (d) Ba₃(PO₄)₂

Although these rules imply that solubility is a simple yes or no question, the real- ity is more complicated than that. Saying that a compound is soluble does not mean that we could dissolve limitless amounts of it in a small beaker of water. If we keep adding more of the solute, eventually we will observe that the added material does not dissolve. When this occurs we have established a saturated solution. The concentration at which a given solution will become saturated depends on the identity of the solute, so it is useful to have a quantitative measure of solubility. The units used to report this number can vary, but one common choice is mass of solute per 100 g of solvent. For example, the solubility of table salt in water at room temperature is 35.7 g NaCl/100 g H₂O. This corresponds to about one-third of a cup of salt in a cup of water. Notice that solubility gives us a ratio, which we could use in the same way as we did the mass density in Example Problem 1.6 on page 20.

One fi nal observation to make about solutions is that many ionic compounds will dissociate into individual ions when they dissolve in water. Thus the salt solution above actually contains Na⁺ and Cl⁻ ions, not NaCl molecules. The availability of freely moving charges allows these solutions to conduct electricity. Any substance that dissolves in water to produce an aqueous solution that conducts electricity is called an electrolyte. Substances whose solutions do not conduct electricity are called nonelectrolytes. We can divide electrolytes further into two groups. Strong electrolytes dissociate completely, so that only individual ions are present in the solution, with virtually no intact molecules. In contrast, weak electrolytes dissociate only

partially; their solutions contain both intact molecules and individual ions in measurable quantities. Figure 3.7 shows the differences between these classes of solutes.

Chemical Equations for Aqueous Reactions

We can describe the process by which a compound dissolves in water by a chemical equation. When a covalently bonded material such as sugar dissolves in water, the molecules remain intact:

C₆H₁₂O₆(s) : C₆H₁₂O₆(aq)

By contrast, when an ionic solid dissolves in water, it breaks into its constituent ions.

This is called a dissociation reaction. The dissolving of sodium chloride discussed above is a common example of this process:

NaCl(s) : Na⁺(aq) + Cl⁻(aq)

In both of these chemical equations, the water molecules are not shown explicitly, although their presence is indicated by the “(aq)” on the product side. This omission refl ects the general tendency of a solvent to dissolve a solute but not to react chemi- cally with it. When water appears in a chemical equation, the reaction involves water as either a reactant or product.

In addition to these fairly simple reactions by which we describe compounds dissolving in water, many important reactions take place in water. The chemical equations we write to describe these reactions can be written in any of three forms; the choice of equation is based mainly on the context in which the equation is used. We’ll Figure 3.7 ❚ The photos show a classroom demonstration in which a pair of copper rods is dipped into different solutions. If the solution conducts electricity, a circuit is closed and the lightbulb lights.

(a) Ordinary sugar (C12H22O11) is a nonelectrolyte, so a solution of sugar dissolved in water does not conduct electricity. (b) Acetic acid (CH₃COOH) is a weak electrolyte. An acetic acid solution contains low concentrations of ions and conducts electricity well enough to light the bulb dimly.

(c) Potassium chromate (K₂CrO₄) is a strong electrolyte. A potassium chromate solution contains higher concentrations of ions and conducts electricity well enough to light the bulb brightly. The illustrations accompanying each photo emphasize the solute species in each solution.

H2O molecule Sugar molecule Hydrogen ion, H⁺

Acetic acid molecule, CH₃COOH

Acetate ion, CH₃COO^–

Potassium ion, K⁺

Chromate ion, CrO₄^2–

(a) (b) (c)

Photos: © Cengage Learning/Charles D. Winters

illustrate these three types of equations for a reaction that is important in the produc- tion of commercial explosives: the synthesis of ammonium nitrate from ammonia and nitric acid.

Ammonium nitrate is a precursor to many important explosives and also has con- siderable explosive potential itself. Because ammonium nitrate is also widely used as a fertilizer, it is readily available for purchase. An unfortunate result of this availability is that it has often been used in bombings, including the 1995 attack on the Murrah Federal Building in Oklahoma City.

In the commercial preparation of ammonium nitrate, pure ammonia (NH₃) in the gas phase is combined with concentrated aqueous nitric acid (HNO₃). This produces a highly concentrated aqueous solution of ammonium nitrate, which is then dried and formed into small pellets (called prills). The molecular equation shows the com- plete formula of each compound involved. Both intact and dissociated compounds are shown, followed by the (aq) to designate an aqueous solution when appropriate:

HNO₃(aq) + NH₃(g) : NH₄NO₃(aq)

Frequently, the compounds involved in aqueous chemistry are ionic, as is true for both HNO₃ and NH₄NO₃ in this example. These ionic compounds dissociate when they dissolve in water, so the actual solution does not contain intact molecules of these compounds. This dissociation is emphasized if we write the total ionic equation rather than the molecular equation. This form emphasizes what is actually present in the reacting mixture by writing dissociated compounds as separated ions in the solution. The total ionic equation for our example reaction from above is thus

H⁺(aq) + NO₃⁻(aq) + NH₃(g) : NH₄⁺(aq) + NO₃⁻(aq)

The H⁺ ion is rather special because it is present in all aqueous solutions of acids.

Strictly speaking, an H⁺ ion would be nothing more than a proton. But as we will see in more detail later, such a bare proton in aqueous solution would actually be tightly sur- rounded by water molecules. Thus the H⁺ ion is frequently written as H₃O⁺ to remind us of this fact. If we write the ion in this way, our total ionic equation becomes

H₃O⁺(aq) + NO₃⁻(aq) + NH₃(g) : NH₄⁺(aq) + NO₃⁻(aq) + H₂O(,) Note that now we must include a molecule of water on the right-hand side to balance the atoms that we added when we switched our representation of the H⁺ ion from H⁺ to H₃O⁺.

Careful inspection of this equation shows that one ion, NO₃⁻(aq), appears on both sides. Thus it is not an active participant in the reaction. Ions that are uninvolved in the chemistry are referred to as spectator ions. In many instances, we are not par- ticularly concerned with these spectator ions and need not include them in the chemical equation. When the spectator ions are left out, the result is a net ionic equation.

Using the same reaction as an example, we would have either H⁺(aq) + NH₃(g) : NH₄⁺(aq) or

H₃O⁺(aq) + NH₃(g) : NH₄⁺(aq) + H₂O(,)

This depends on whether we choose to write H⁺ or H₃O⁺ to represent the cation from the nitric acid. Because we have removed anions from both sides of the equation, we now have a net charge of 1+ on each side. This is one way to recognize that you are looking at a net ionic equation: there is often a nonzero net charge on each side of the equation. If the equation is properly balanced, though, the net charge on each side must be the same. This rule refl ects the fact that electrical charge must be conserved.

At this point, you are probably wondering, “Which of these equations is the correct one?” The answer is that all of these forms provide valid descriptions of the reaction, so none is intrinsically “better” than the others. In this particular case, we would probably not choose the net ionic equation because the identity of the anion (NO₃⁻)

The process of removing the spectator ions from the equation is similar to what you would do if you encountered an algebraic equation that had the same term on both sides.

is important if our aim is to produce solid ammonium nitrate for the explosives industry. But in many other cases, we might have little or no interest in the identity of a spectator ion, so the simplicity of the net ionic equation might make it attractive. The total ionic equation is more cumbersome and is used less often.

Acid–Base Reactions

Two especially important categories of aqueous solutions are acids and bases. Exam- ples of them are easy to fi nd in our everyday lives as well as in the chemical industry.

For our present purposes, we will defi ne an acid as any substance that dissolves in water to produce H⁺ (or H₃O⁺) ions and a base as any substance that dissolves in water to produce OH⁻ ions. Table 3.2 lists some common acids and bases. Like other solutes, acids and bases can be either strong or weak electrolytes. Strong acids or bases dissociate completely in water, so that the resulting solution contains essentially no intact solute molecules. We can write the following chemical equations for the dissociation of HCl and NaOH:

Strong acid: HCl(g) + H₂O(ℓ) : H₃O⁺(aq) + Cl⁻(aq) Strong base: NaOH(s) : Na⁺(aq) + OH⁻(aq)

The only common strong acids and bases are those shown in Table 3.2.

Weak acids or bases undergo only partial ionization, so that their solutions contain intact molecules as well as dissociated ions. When we write equations for the dissolution of weak electrolytes like these, we use a “two-way” arrow that emphasizes that the reaction does not proceed completely from left to right. For common weak acids, it is relatively easy to write the needed ionization equation. Many weak acids contain the !COOH functional group, and the H atom from that group tends to be lost in solution. Acetic acid is a good example.

Weak acid: CH₃COOH(aq) + H₂O(ℓ) N H₃O⁺(aq) + CH₃COO⁻(aq) Unlike H⁺ ions, hydroxide ions are not

represented as having combined with water from the solvent, so water doesn’t appear in the equation for a strong base.

Unlike H⁺ ions, hydroxide ions are not represented as having combined with water from the solvent, so water doesn’t appear in the equation for a strong base.

We will revisit the concepts of acids and bases in Chapter 12 and provide additional detail at that time.

Table

❚

^3.2

Strong and weak acids and bases

Strong Acids Strong Bases

HCl Hydrochloric acid LiOH Lithium hydroxide

HNO3 Nitric acid NaOH Sodium hydroxide

H2SO4 Sulfuric acid KOH Potassium hydroxide

HClO₄ Perchloric acid Ca(OH)₂ Calcium hydroxide

HBr Hydrobromic acid Ba(OH)₂ Barium hydroxide

HI Hydriodic acid Sr(OH)2 Strontium hydroxide

Weak Acids Weak Bases

H3PO4 Phosphoric acid NH3 Ammonia

HF Hydrofl uoric acid CH3NH2 Methylamine

CH3COOH Acetic acid

HCN Hydrocyanic acid

Note: All common strong acids and bases are shown, but only representative examples of weak acids and bases are listed.

For weak bases, the situation is slightly less obvious. Unlike the strong bases listed in Table 3.2, most weak bases do not contain !OH groups. So it may not be obvious that dissolving them will produce hydroxide ions in solution (i.e., that they are bases at all). The most common weak base is ammonia (NH₃), which reacts in water accord- ing to the following equation:

Weak base: NH₃(aq) + H₂O(ℓ) N NH₄⁺(aq) + OH⁻(aq)

Again we use the two-way arrow to show that intact NH₃ molecules as well as NH₄⁺ and OH⁻ ions will be present in the solution. Many other weak bases are amines.

They can be thought of as derivatives of ammonia in which one or more of the H atoms have been replaced by methyl groups or longer hydrocarbon chains.

Because many acidic and basic solutions occur in nature, observations about acids and bases date back hundreds of years. One of the most important observations is that a solution cannot be both acidic and basic at the same time. Mixing an acid and a base leads to a reaction known as neutralization, in which the resulting solution is neither acidic nor basic. To understand the origins of this neutralization, we might start with the defi nitions of acid and base. Acid solutions contain H₃O⁺ ions, and bases contain OH⁻ ions. So for a solution to be both an acid and a base simultaneously, it would need to contain both of these species. Looking at those ions, it should be easy to see why this is not feasible. Hydronium ions and hydroxide ions combine readily to form water:

H₃O⁺(aq) + OH⁻(aq) : 2 H₂O(,)

This reaction will take place whenever an acid and a base are combined and will always prevent a solution from being both acidic and basic simultaneously.

E X A M P L E P RO B L E M 3 . 3

When aqueous solutions of acetic acid and potassium hydroxide are combined, a neutralization reaction will occur. Write molecular, total ionic, and net ionic equations for this process.

Strategy This is the reaction of a weak acid with a strong base. The reaction between an acid and a base always produces water as a product, along with an ionic compound formed from the remaining ions. (This second product is often called a salt.) We can begin by using this idea to write the molecular equation. To generate the ionic equations, then, we account for the dissociation of strong electrolytes into their constituent ions.

Solution A hydrogen atom from the acid and the hydroxide ion from the base will produce water. The remaining ions will be an acetate anion (CH₃COO⁻ from the acid) and a potassium ion (K⁺, from the base). Combining these gives us potassium acetate, KCH₃COO. This lets us write the molecular equation:

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