Classifying Chemical Reactions

Learning Objectives

By the end of this section, you will be able to:

  • Define three common types of chemical reactions (precipitation, acid-base, and oxidation-reduction)
  • Classify chemical reactions as one of these three types given appropriate descriptions or chemical equations
  • Identify common acids and bases
  • Predict the solubility of common inorganic compounds by using solubility rules
  • Compute the oxidation states for elements in compounds

Humans interact with one another in various and complex ways, and we classify these interactions according to common patterns of behavior. When two humans exchange information, we say they are communicating. When they exchange blows with their fists or feet, we say they are fighting. Faced with a wide range of varied interactions between chemical substances, scientists have likewise found it convenient (or even necessary) to classify chemical interactions by identifying common patterns of reactivity. This module will provide an introduction to three of the most prevalent types of chemical reactions: precipitation, acid-base, and oxidation-reduction.

Precipitation Reactions and Solubility Rules

A precipitation reaction is one in which dissolved substances react to form one (or more) solid products. Many reactions of this type involve the exchange of ions between ionic compounds in aqueous solution and are sometimes referred to as double displacement, double replacement, or metathesis reactions. These reactions are common in nature and are responsible for the formation of coral reefs in ocean waters and kidney stones in animals. They are used widely in industry for production of a number of commodity and specialty chemicals. Precipitation reactions also play a central role in many chemical analysis techniques, including spot tests used to identify metal ions and gravimetric methods for determining the composition of matter (see the last module of this chapter).

The extent to which a substance may be dissolved in water, or any solvent, is quantitatively expressed as its solubility, defined as the maximum concentration of a substance that can be achieved under specified conditions. Substances with relatively large solubilities are said to be soluble. A substance will precipitate when solution conditions are such that its concentration exceeds its solubility. Substances with relatively low solubilities are said to be insoluble, and these are the substances that readily precipitate from solution. More information on these important concepts is provided in the text chapter on solutions. For purposes of predicting the identities of solids formed by precipitation reactions, one may simply refer to patterns of solubility that have been observed for many ionic compounds.

Solubilities of Common Ionic Compounds in Water

Soluble compounds contain

  • group 1 metal cations (Li+, Na+, K+, Rb+, and Cs+) and ammonium ion NH4+
  • the halide ions (Cl, Br, and I)
  • the acetate (C2H3O2), bicarbonate (HCO3) , nitrate (NO3), and chlorate (ClO3) ions
  • the sulfate (SO4) ion

Exceptions to these solubility rules include

  • halides of Ag+, Hg22+, and Pb2+
  • sulfates of Ag+, Ba2+, Ca2+, Hg22+, Pb2+, and Sr2+

Insoluble compounds contain

  • carbonate (CO32−), chromate (CrO32−), phosphate (PO43−) and sulfide (S2–) ions
  • hydroxide ion (OH)

Exceptions to these insolubility rules include

  • compounds of these anions with group 1 metal cations and ammonium ion
  • hydroxides of group 1 metal cations and Ba2+

A vivid example of precipitation is observed when solutions of potassium iodide and lead nitrate are mixed, resulting in the formation of solid lead iodide:


This observation is consistent with the solubility guidelines: The only insoluble compound among all those involved is lead iodide, one of the exceptions to the general solubility of iodide salts.

The net ionic equation representing this reaction is:


A photograph is shown of a yellow green opaque substance swirled through a clear, colorless liquid in a test tube. Figure 1. A precipitate of PbI2 forms when solutions containing Pb2+ and I are mixed. (credit: Der Kreole/Wikimedia Commons)

Lead iodide is a bright yellow solid that was formerly used as an artist’s pigment known as iodine yellow (Figure 1). The properties of pure PbI2 crystals make them useful for fabrication of X-ray and gamma ray detectors.

The solubility guidelines discussed above may be used to predict whether a precipitation reaction will occur when solutions of soluble ionic compounds are mixed together. One merely needs to identify all the ions present in the solution and then consider if possible cation/anion pairing could result in an insoluble compound.

For example, mixing solutions of silver nitrate and sodium fluoride will yield a solution containing Ag+, NO, Na+, and F ions. Aside from the two ionic compounds originally present in the solutions, AgNO3 and NaF, two additional ionic compounds may be derived from this collection of ions: NaNO3 and AgF. The solubility guidelines indicate all nitrate salts are soluble but that AgF is one of the exceptions to the general solubility of fluoride salts. A precipitation reaction, therefore, is predicted to occur, as described by the following equations:

NaF(aq)+AgNO3(aq)AgF(s)+NaNO3(aq)(molecular)Ag+(aq)+F(aq)AgF(s)(net ionic)\begin{array}{l}\text{NaF(}aq\text{)}+{\text{AgNO}}_{3}\text{(}aq\text{)}\rightarrow\text{AgF(}s\text{)}+\text{NaN}{\text{O}}_{3}\text{(}aq\text{)}\text{(molecular)}\\ {\text{Ag}}^{\text{+}}\text{(}aq\text{)}+{\text{F}}^{-}\text{(}aq\text{)}\rightarrow\text{AgF(}s\text{)}\text{(net ionic)}\end{array}

Example 1: Predicting Precipitation Reactions

Predict the result of mixing reasonably concentrated solutions of the following ionic compounds. If precipitation is expected, write a balanced net ionic equation for the reaction.

  1. potassium sulfate and barium nitrate
  2. lithium chloride and silver acetate
  3. lead nitrate and ammonium carbonate

Check Your Learning

Which solution could be used to precipitate the barium ion, Ba2+, in a water sample: sodium chloride, sodium hydroxide, or sodium sulfate? What is the formula for the expected precipitate?

Acid-Base Reactions

An acid-base reaction is one in which a hydrogen ion, H+, is transferred from one chemical species to another. Such reactions are of central importance to numerous natural and technological processes, ranging from the chemical transformations that take place within cells and the lakes and oceans, to the industrial-scale production of fertilizers, pharmaceuticals, and other substances essential to society. The subject of acid-base chemistry, therefore, is worthy of thorough discussion, and a full chapter is devoted to this topic later in the text.

For purposes of this brief introduction, we will consider only the more common types of acid-base reactions that take place in aqueous solutions. In this context, an acid is a substance that will dissolve in water to yield hydronium ions, H3O+. As an example, consider the equation shown here:


The process represented by this equation confirms that hydrogen chloride is an acid. When dissolved in water, H3O+ ions are produced by a chemical reaction in which H+ ions are transferred from HCl molecules to H2O molecules (Figure 2).

This figure shows two flasks, labeled a and b. The flasks are both sealed with stoppers and are nearly three-quarters full of a liquid. Flask a is labeled H C l followed by g in parentheses. In the liquid there are approximately twenty space-filling molecular models composed of one red sphere and two smaller attached white spheres. The label H subscript 2 O followed by a q in parentheses is connected with a line to one of these models. In the space above the liquid in the flask, four space filling molecular models composed of one larger green sphere to which a smaller white sphere is bonded are shown. To one of these models, the label H C l followed by g in parentheses is attached with a line segment. An arrow is drawn from the space above the liquid pointing down into the liquid below. Flask b is labeled H subscript 3 O superscript positive sign followed by a q in parentheses. This is followed by a plus sign and C l superscript negative sign which is also followed by a q in parentheses. In this flask, no molecules are shown in the open space above the liquid. A label, C l superscript negative sign followed by a q in parentheses, is connected with a line segment to a green sphere. This sphere is surrounded by four molecules composed each of one red sphere and two white smaller spheres. A few of these same molecules appear separate from the green spheres in the liquid. A line segment connects one of them to the label H subscript 2 O which is followed by l in parentheses. There are a few molecules formed from one central larger red sphere to which three smaller white spheres are bonded. A line segment is drawn from one of these to the label H subscript 3 O superscript positive sign, followed by a q in parentheses. Figure 2. When hydrogen chloride gas dissolves in water, (a) it reacts as an acid, transferring protons to water molecules to yield (b) hydronium ions (and solvated chloride ions).

This figure contains two images, each with an associated structural formula provided in the lower left corner of the image. The first image is a photograph of a variety of thinly sliced, circular cross sections of citrus fruits ranging in color for green to yellow, to orange and reddish-orange. The slices are closely packed on a white background. The structural formula with this picture shows a central chain of five C atoms. The leftmost C atom has an O atom double bonded above and to the left and a singly bonded O atom below and to the left. This single bonded O atom has an H atom indicated in red on its left side which is highlighted in pink. The second C atom moving to the right has H atoms bonded above and below. The third C atom has a single bonded O atom above which has an H atom on its right. This third C atom has a C atom bonded below it which has an O atom double bonded below and to the left and a singly bonded O atom below and to the right. An H atom appears in red and is highlighted in pink to the right of the singly bonded O atom. The fourth C atom has H atoms bonded above and below. The fifth C atom is at the right end of the structure. It has an O atom double bonded above and to the right and a singly bonded O atom below and to the right. This single bonded O atom has a red H atom on its right side which is highlighted in pink. The second image is a photograph of bottles of vinegar. The bottles are labeled, “Balsamic Vinegar,” and appear to be clear and colorless. The liquid in this bottle appears to be brown. The structural formula that appears with this image shows a chain of two C atoms. The leftmost C atom has H atoms bonded above, below, and to the left. The C atom on the right has a doubly bonded O atom above and to the right and a singly bonded O atom below and to the right. This O atom has an H atom bonded to its right which is highlighted in pink. Figure 3. (a) Fruits such as oranges, lemons, and grapefruit contain the weak acid citric acid. (b) Vinegars contain the weak acid acetic acid. The hydrogen atoms that may be transferred during an acid-base reaction are highlighted in the inset molecular structures. (credit a: modification of work by Scott Bauer; credit b: modification of work by Brücke-Osteuropa/Wikimedia Commons)

The nature of HCl is such that its reaction with water as just described is essentially 100% efficient: Virtually every HCl molecule that dissolves in water will undergo this reaction. Acids that completely react in this fashion are called strong acids, and HCl is one among just a handful of common acid compounds that are classified as strong (Table 1).

A far greater number of compounds behave as weak acids and only partially react with water, leaving a large majority of dissolved molecules in their original form and generating a relatively small amount of hydronium ions. Weak acids are commonly encountered in nature, being the substances partly responsible for the tangy taste of citrus fruits, the stinging sensation of insect bites, and the unpleasant smells associated with body odor. A familiar example of a weak acid is acetic acid, the main ingredient in food vinegars:


When dissolved in water under typical conditions, only about 1% of acetic acid molecules are present in the ionized form,
(Figure 3). (The use of a double-arrow in the equation above denotes the partial reaction aspect of this process, a concept addressed fully in the chapters on chemical equilibrium.)

Table 1. Common Strong Acids
Compound Formula Name in Aqueous Solution
HBr hydrobromic acid
HCl hydrochloric acid
HI hydroiodic acid
HNO3 nitric acid
HClO4 perchloric acid
H2SO4 sulfuric acid
A base is a substance that will dissolve in water to yield hydroxide ions, OH. The most common bases are ionic compounds composed of alkali or alkaline earth metal cations (groups 1 and 2) combined with the hydroxide ion—for example, NaOH and Ca(OH)2. When these compounds dissolve in water, hydroxide ions are released directly into the solution. For example, KOH and Ba(OH)2 dissolve in water and dissociate completely to produce cations (K+ and Ba2+, respectively) and hydroxide ions, OH. These bases, along with other hydroxides that completely dissociate in water, are considered strong bases.

Consider as an example the dissolution of lye (sodium hydroxide) in water:


This equation confirms that sodium hydroxide is a base. When dissolved in water, NaOH dissociates to yield Na+ and OH ions. This is also true for any other ionic compound containing hydroxide ions. Since the dissociation process is essentially complete when ionic compounds dissolve in water under typical conditions, NaOH and other ionic hydroxides are all classified as strong bases.

Unlike ionic hydroxides, some compounds produce hydroxide ions when dissolved by chemically reacting with water molecules. In all cases, these compounds react only partially and so are classified as weak bases. These types of compounds are also abundant in nature and important commodities in various technologies. For example, global production of the weak base ammonia is typically well over 100 metric tons annually, being widely used as an agricultural fertilizer, a raw material for chemical synthesis of other compounds, and an active ingredient in household cleaners (Figure 4). When dissolved in water, ammonia reacts partially to yield hydroxide ions, as shown here:


This is, by definition, an acid-base reaction, in this case involving the transfer of H+ ions from water molecules to ammonia molecules. Under typical conditions, only about 1% of the dissolved ammonia is present as

This photograph shows a large agricultural tractor in a field pulling a field sprayer and a large, white cylindrical tank which is labeled “Caution Ammonia.” Figure 4. Ammonia is a weak base used in a variety of applications. (a) Pure ammonia is commonly applied as an agricultural fertilizer. (b) Dilute solutions of ammonia are effective household cleansers. (credit a: modification of work by National Resources Conservation Service; credit b: modification of work by pat00139)

The chemical reactions described in which acids and bases dissolved in water produce hydronium and hydroxide ions, respectively, are, by definition, acid-base reactions. In these reactions, water serves as both a solvent and a reactant. A neutralization reaction is a specific type of acid-base reaction in which the reactants are an acid and a base, the products are often a salt and water, and neither reactant is the water itself:


To illustrate a neutralization reaction, consider what happens when a typical antacid such as milk of magnesia (an aqueous suspension of solid Mg(OH)2) is ingested to ease symptoms associated with excess stomach acid (HCl):


Note that in addition to water, this reaction produces a salt, magnesium chloride.

Example 2: Writing Equations for Acid-Base Reactions

Write balanced chemical equations for the acid-base reactions described here:

  1. the weak acid hydrogen hypochlorite reacts with water
  2. a solution of barium hydroxide is neutralized with a solution of nitric acid

Check Your Learning

Write the net ionic equation representing the neutralization of any strong acid with an ionic hydroxide. (Hint: Consider the ions produced when a strong acid is dissolved in water.)

Explore the microscopic view of strong and weak acids and bases at the PhET Acid-Base Simulator.

Oxidation-Reduction Reactions

Earth’s atmosphere contains about 20% molecular oxygen, O2, a chemically reactive gas that plays an essential role in the metabolism of aerobic organisms and in many environmental processes that shape the world. The term oxidation was originally used to describe chemical reactions involving O2, but its meaning has evolved to refer to a broad and important reaction class known as oxidation-reduction (redox) reactions. A few examples of such reactions will be used to develop a clear picture of this classification.

Some redox reactions involve the transfer of electrons between reactant species to yield ionic products, such as the reaction between sodium and chlorine to yield sodium chloride:


It is helpful to view the process with regard to each individual reactant, that is, to represent the fate of each reactant in the form of an equation called a half-reaction:

2Na(s)2Na+(s)+2eCl2(g)+2e2Cl(s)\begin{array}{l}2\text{Na(}s\text{)}\rightarrow 2{\text{Na}}^{\text{+}}\text{(}s\text{)}+2{\text{e}}^{-}\\ {\text{Cl}}_{2}\text{(}g\text{)}+2{\text{e}}^{-}\rightarrow 2{\text{Cl}}^{-}\text{(}s\text{)}\end{array}

These equations show that Na atoms lose electrons while Cl atoms (in the Cl2 molecule) gain electrons, the “s” subscripts for the resulting ions signifying they are present in the form of a solid ionic compound. For redox reactions of this sort, the loss and gain of electrons define the complementary processes that occur:

oxidation=loss of electronsreduction=gain of electrons\begin{array}{lll}\qquad \mathbf{\text{oxidation}}& =& \text{loss of electrons}\qquad \\ \qquad \mathbf{\text{reduction}}& =& \text{gain of electrons}\qquad \end{array}

In this reaction, then, sodium is oxidized and chlorine is undergoes reduction. Viewed from a more active perspective, sodium functions as a reducing agent (reductant), since it provides electrons to (or reduces) chlorine. Likewise, chlorine functions as an oxidizing agent (oxidant), as it effectively removes electrons from (oxidizes) sodium.

reducing agent=species that is oxidizedoxidizing agent=species that is reduced\begin{array}{lll}\qquad \mathbf{\text{reducing agent}}& =& \text{species that is oxidized}\qquad \\ \qquad \mathbf{\text{oxidizing agent}}& =& \text{species that is reduced}\qquad \end{array}

Some redox processes, however, do not involve the transfer of electrons. Consider, for example, a reaction similar to the one yielding NaCl:

H2(g)+Cl2(g)2HCl(g){\text{H}}_{2}\text{(}g\text{)}+{\text{Cl}}_{2}\text{(}g\text{)}\rightarrow 2\text{HCl(}g\text{)}

The product of this reaction is a covalent compound, so transfer of electrons in the explicit sense is not involved. To clarify the similarity of this reaction to the previous one and permit an unambiguous definition of redox reactions, a property called oxidation number has been defined. The oxidation number (or oxidation state) of an element in a compound is the charge its atoms would possess if the compound was ionic. The following guidelines are used to assign oxidation numbers to each element in a molecule or ion.

  1. The oxidation number of an atom in an elemental substance is zero.
  2. The oxidation number of a monatomic ion is equal to the ion’s charge.
  3. Oxidation numbers for common nonmetals are usually assigned as follows:

    • Hydrogen: +1 when combined with nonmetals, –1 when combined with metals
    • Oxygen: –2 in most compounds, sometimes –1 (so-called peroxides,
      very rarely
      (so-called superoxides,
      positive values when combined with F (values vary)
    • Halogens: –1 for F always, –1 for other halogens except when combined with oxygen or other halogens (positive oxidation numbers in these cases, varying values)

  4. The sum of oxidation numbers for all atoms in a molecule or polyatomic ion equals the charge on the molecule or ion.

Note: The proper convention for reporting charge is to write the number first, followed by the sign (e.g., 2+), while oxidation number is written with the reversed sequence, sign followed by number (e.g., +2). This convention aims to emphasize the distinction between these two related properties.

Example 3: Assigning Oxidation Numbers

Follow the guidelines in this section of the text to assign oxidation numbers to all the elements in the following species:

  1. H2S
  2. SO32{\text{SO}}_{3}{}^{\text{2}-}
  3. Na2SO4

Check Your Learning

Assign oxidation states to the elements whose atoms are underlined in each of the following compounds or ions:

  1. KNO3
  2. AlH3
  3. NH4+\underline{\text{N}}{\text{H}}_{4}{}^{+}
  4. H2PO4{\text{H}}_{2}\underline{\text{P}}{\text{O}}_{4}{}^{-}

Using the oxidation number concept, an all-inclusive definition of redox reaction has been established. Oxidation-reduction (redox) reactions are those in which one or more elements involved undergo a change in oxidation number. (While the vast majority of redox reactions involve changes in oxidation number for two or more elements, a few interesting exceptions to this rule do exist, as in Example 4.) Definitions for the complementary processes of this reaction class are correspondingly revised as shown here:

oxidation=increase in oxidation numberreduction=decrease in oxidation number\begin{array}{lll}\qquad \mathbf{\text{oxidation}}& =& \text{increase in oxidation number}\qquad \\ \qquad \mathbf{\text{reduction}}& =& \text{decrease in oxidation number}\qquad \end{array}

Returning to the reactions used to introduce this topic, they may now both be identified as redox processes. In the reaction between sodium and chlorine to yield sodium chloride, sodium is oxidized (its oxidation number increases from 0 in Na to +1 in NaCl) and chlorine is reduced (its oxidation number decreases from 0 in Cl2 to –1 in NaCl). In the reaction between molecular hydrogen and chlorine, hydrogen is oxidized (its oxidation number increases from 0 in H2 to +1 in HCl) and chlorine is reduced (its oxidation number decreases from 0 in Cl2 to –1 in HCl).

Several subclasses of redox reactions are recognized, including combustion reactions in which the reductant (also called a fuel) and oxidant (often, but not necessarily, molecular oxygen) react vigorously and produce significant amounts of heat, and often light, in the form of a flame. Solid rocket-fuel reactions are combustion processes. A typical propellant reaction in which solid aluminum is oxidized by ammonium perchlorate is represented by this equation:

10Al(s)+6NH4ClO4(s)4Al2O3(s)+2AlCl3(s)+12H2O(g)+3N2(g)10\text{Al(}s\text{)}+6{\text{NH}}_{4}{\text{ClO}}_{4}\text{(}s\text{)}\rightarrow 4{\text{Al}}_{2}{\text{O}}_{3}\text{(}s\text{)}+2{\text{AlCl}}_{3}\text{(}s\text{)}+12{\text{H}}_{2}\text{O(}g\text{)}+3{\text{N}}_{2}\text{(}g\text{)}

Watch a brief video showing the test firing of a small-scale, prototype, hybrid rocket engine planned for use in the new Space Launch System being developed by NASA. The first engines firing at 3 s (green flame) use a liquid fuel/oxidant mixture, and the second, more powerful engines firing at 4 s (yellow flame) use a solid mixture.

Single-displacement (replacement) reactions are redox reactions in which an ion in solution is displaced (or replaced) via the oxidation of a metallic element. One common example of this type of reaction is the acid oxidation of certain metals:


Metallic elements may also be oxidized by solutions of other metal salts; for example:


This reaction may be observed by placing copper wire in a solution containing a dissolved silver salt. Silver ions in solution are reduced to elemental silver at the surface of the copper wire, and the resulting Cu2+ ions dissolve in the solution to yield a characteristic blue color (Figure 5).

This figure contains three photographs. In a, a coiled copper wire is shown beside a test tube filled with a clear, colorless liquid. In b, the wire has been inserted into the test tube with the clear, colorless liquid. In c, the test tube contains a light blue liquid and the coiled wire appears to have a fuzzy silver gray coating. Figure 5. (a) A copper wire is shown next to a solution containing silver(I) ions. (b) Displacement of dissolved silver ions by copper ions results in (c) accumulation of gray-colored silver metal on the wire and development of a blue color in the solution, due to dissolved copper ions. (credit: modification of work by Mark Ott)

Example 4: Describing Redox Reactions

Identify which equations represent redox reactions, providing a name for the reaction if appropriate. For those reactions identified as redox, name the oxidant and reductant.

  1. ZnCO3(s)ZnO(s)+CO2(g){\text{ZnCO}}_{3}\text{(}s\text{)}\rightarrow\text{ZnO(}s\text{)}+{\text{CO}}_{2}\text{(}g\text{)}
  2. 2Ga(l)+3Br2(l)2GaBr3(s)2\text{Ga}\text{(}l\text{)}+3{\text{Br}}_{2}\text{(}l\text{)}\rightarrow 2{\text{GaBr}}_{3}\text{(}s\text{)}
  3. 2H2O2(aq)2H2O(l)+O2(g)2{\text{H}}_{2}{\text{O}}_{2}\text{(}aq\text{)}\rightarrow 2{\text{H}}_{2}\text{O}\text{(}l\text{)}+{\text{O}}_{2}\text{(}g\text{)}
  4. BaCl2(aq)+K2SO4(aq)BaSO4(s)+2KCl(aq){\text{BaCl}}_{2}\text{(}aq\text{)}+{\text{K}}_{2}{\text{SO}}_{4}\text{(}aq\text{)}\rightarrow{\text{BaSO}}_{4}\text{(}s\text{)}+2\text{KCl}\text{(}aq\text{)}
  5. C2H4(g)+3O2(g)2CO2(g)+2H2O(l){\text{C}}_{2}{\text{H}}_{4}\text{(}g\text{)}+3{\text{O}}_{2}\text{(}g\text{)}\rightarrow 2{\text{CO}}_{2}\text{(}g\text{)}+2{\text{H}}_{2}\text{O}\text{(}l\text{)}

Check Your Learning

This equation describes the production of tin(II) chloride:

Is this a redox reaction? If so, provide a more specific name for the reaction if appropriate, and identify the oxidant and reductant.

Balancing Redox Reactions via the Half-Reaction Method

Redox reactions that take place in aqueous media often involve water, hydronium ions, and hydroxide ions as reactants or products. Although these species are not oxidized or reduced, they do participate in chemical change in other ways (e.g., by providing the elements required to form oxyanions). Equations representing these reactions are sometimes very difficult to balance by inspection, so systematic approaches have been developed to assist in the process. One very useful approach is to use the method of half-reactions, which involves the following steps:

  1. Write the two half-reactions representing the redox process.
  2. Balance all elements except oxygen and hydrogen.
  3. Balance oxygen atoms by adding H2O molecules.
  4. Balance hydrogen atoms by adding H+ ions.
  5. Balance charge by adding electrons.
  6. If necessary, multiply each half-reaction’s coefficients by the smallest possible integers to yield equal numbers of electrons in each.
  7. Add the balanced half-reactions together and simplify by removing species that appear on both sides of the equation.
  8. For reactions occurring in basic media (excess hydroxide ions), carry out these additional steps:

    • Add OH ions to both sides of the equation in numbers equal to the number of H+ ions.
    • On the side of the equation containing both H+ and OH ions, combine these ions to yield water molecules.
    • Simplify the equation by removing any redundant water molecules.

  9. Finally, check to see that both the number of atoms and the total charges2 are balanced.

Example 5: Balancing Redox Reactions in Acidic Solution

Write a balanced equation for the reaction between dichromate ion and iron(II) to yield iron(III) and chromium(III) in acidic solution.


Check Your Learning

In acidic solution, hydrogen peroxide reacts with Fe2+ to produce Fe3+ and H2O. Write a balanced equation for this reaction.

Key Concepts and Summary

Chemical reactions are classified according to similar patterns of behavior. A large number of important reactions are included in three categories: precipitation, acid-base, and oxidation-reduction (redox). Precipitation reactions involve the formation of one or more insoluble products. Acid-base reactions involve the transfer of hydrogen ions between reactants. Redox reactions involve a change in oxidation number for one or more reactant elements. Writing balanced equations for some redox reactions that occur in aqueous solutions is simplified by using a systematic approach called the half-reaction method.


  1. Use the following equations to answer the next four questions:

    • H2O(s)H2O(l){\text{H}}_{2}\text{O(}s\text{)}\rightarrow{\text{H}}_{2}\text{O(}l\text{)}
    • Na+(aq)+Cl(aq)Ag+(aq)+NO3(aq)AgCl(s)+Na+(aq)+NO3(aq){\text{Na}}^{+}\text{(}aq\text{)}+{\text{Cl}}^{-}\text{(}aq\text{)}{\text{Ag}}^{+}\text{(}aq\text{)}+{\text{NO}}_{3}{}^{-}\text{(}aq\text{)}\rightarrow\text{AgCl(}s\text{)}+{\text{Na}}^{+}\text{(}aq\text{)}+{\text{NO}}_{3}{}^{-}\text{(}aq\text{)}
    • CH3OH(g)+O2(g)CO2(g)+H2O(g){\text{CH}}_{3}\text{OH(}g\text{)}+{\text{O}}_{2}\text{(}g\text{)}\rightarrow{\text{CO}}_{2}\text{(}g\text{)}+{\text{H}}_{2}\text{O(}g\text{)}
    • 2H2O(l)2H2(g)+O2(g)2{\text{H}}_{2}\text{O(}l\text{)}\rightarrow 2{\text{H}}_{2}\text{(}g\text{)}+{\text{O}}_{2}\text{(}g\text{)}
    • H+(aq)+OH(aq)H2O(l){\text{H}}^{\text{+}}\text{(}aq\text{)}+{\text{OH}}^{-}\text{(}aq\text{)}\rightarrow{\text{H}}_{2}\text{O(}l\text{)}

    1. Which equation describes a physical change?
    2. Which equation identifies the reactants and products of a combustion reaction?
    3. Which equation is not balanced?
    4. Which is a net ionic equation?

  2. Indicate what type, or types, of reaction each of the following represents:

    1. Ca(s)+Br2(l)CaBr2(s)\text{Ca(}s\text{)}+{\text{Br}}_{2}\text{(}l\text{)}\rightarrow{\text{CaBr}}_{2}\text{(}s\text{)}
    2. Ca(OH)2(aq)+2HBr(aq)CaBr2(aq)+2H2O(l)\text{Ca}{\text{(OH)}}_{2}\text{(}aq\text{)}+2\text{HBr(}aq\text{)}\rightarrow{\text{CaBr}}_{2}\text{(}aq\text{)}+2{\text{H}}_{2}\text{O(}l\text{)}
    3. C6H12(l)+9O2(g)6CO2(g)+6H2O(g){\text{C}}_{6}{\text{H}}_{12}\text{(}l\text{)}+9{\text{O}}_{2}\text{(}g\text{)}\rightarrow 6{\text{CO}}_{2}\text{(}g\text{)}+6{\text{H}}_{2}\text{O(}g\text{)}

  3. Indicate what type, or types, of reaction each of the following represents:

    1. H2O(g)+C(s)CO(g)+H2(g){\text{H}}_{2}\text{O(}g\text{)}+\text{C(}s\text{)}\rightarrow\text{CO(}g\text{)}+{\text{H}}_{2}\text{(}g\text{)}
    2. 2KClO3(s)2KCl(s)+3O2(g){\text{2KClO}}_{3}\text{(}s\text{)}\rightarrow 2\text{KCl(}s\text{)}+3{\text{O}}_{2}\text{(}g\text{)}
    3. Al(OH)3(aq)+3HCl(aq)AlBr3(aq)+3H2O(l)\text{Al}{\text{(OH)}}_{3}\text{(}aq\text{)}+3\text{HCl(}aq\text{)}\rightarrow{\text{AlBr}}_{3}\text{(}aq\text{)}+3{\text{H}}_{2}\text{O}\text{(}l\text{)}
    4. Pb(NO3)2(aq)+H2SO4(aq)PbSO4(s)+2HNO3(aq)\text{Pb}{\text{(}{\text{NO}}_{3}\text{)}}_{2}\text{(}aq\text{)}+{\text{H}}_{2}{\text{SO}}_{4}\text{(}aq\text{)}\rightarrow{\text{PbSO}}_{4}\text{(}s\text{)}+2{\text{HNO}}_{3}\text{(}aq\text{)}

  4. Silver can be separated from gold because silver dissolves in nitric acid while gold does not. Is the dissolution of silver in nitric acid an acid-base reaction or an oxidation-reduction reaction? Explain your answer.
  5. Determine the oxidation states of the elements in the following compounds:

    1. NaI
    2. GdCl3
    3. LiNO3
    4. H2Se
    5. Mg2Si
    6. RbO2, rubidium superoxide
    7. HF

  6. Determine the oxidation states of the elements in the compounds listed. None of the oxygen-containing compounds are peroxides or superoxides.

    1. H3PO4
    2. Al(OH)3
    3. SeO2
    4. KNO2
    5. In2S3
    6. P4O6

  7. Determine the oxidation states of the elements in the compounds listed. None of the oxygen-containing compounds are peroxides or superoxides.

    1. H2SO4
    2. Ca(OH)2
    3. BrOH
    4. ClNO2
    5. TiCl4
    6. NaH

  8. Classify the following as acid-base reactions or oxidation-reduction reactions:

    1. Na2S(aq)+2HCl(aq)2NaCl(aq)+H2S(g){\text{Na}}_{2}\text{S(}aq\text{)}+2\text{HCl(}aq\text{)}\rightarrow 2\text{NaCl(}aq\text{)}+{\text{H}}_{2}\text{S(}g\text{)}
    2. 2Na(s)+2HCl(aq)2NaCl(aq)+H2(g)2\text{Na(}s\text{)}+2\text{HCl(}aq\text{)}\rightarrow 2\text{NaCl(}aq\text{)}+{\text{H}}_{2}\text{(}g\text{)}
    3. Mg(s)+Cl2(g)MgCl2(s)\text{Mg(}s\text{)}+{\text{Cl}}_{2}\text{(}g\text{)}\rightarrow{\text{MgCl}}_{2}\text{(}s\text{)}
    4. MgO(s)+2HCl(aq)MgCl2(aq)+H2O(l)\text{MgO(}s\text{)}+2\text{HCl(}aq\text{)}\rightarrow{\text{MgCl}}_{2}\text{(}aq\text{)}+{\text{H}}_{2}\text{O(}l\text{)}
    5. K3P(s)+2O2(g)K3PO4(s){\text{K}}_{3}\text{P(}s\text{)}+2{\text{O}}_{2}\text{(}g\text{)}\rightarrow{\text{K}}_{3}{\text{PO}}_{4}\text{(}s\text{)}
    6. 3KOH(aq)+H3PO4(aq)K3PO4(aq)+3H2O(l)3\text{KOH(}aq\text{)}+{\text{H}}_{3}{\text{PO}}_{4}\text{(}aq\text{)}\rightarrow{\text{K}}_{3}{\text{PO}}_{4}\text{(}aq\text{)}+3{\text{H}}_{2}\text{O(}l\text{)}

  9. Identify the atoms that are oxidized and reduced, the change in oxidation state for each, and the oxidizing and reducing agents in each of the following equations:

    1. Mg(s)+NiCl2(aq)MgCl2(aq)+Ni(s)\text{Mg(}s\text{)}+{\text{NiCl}}_{2}\text{(}aq\text{)}\rightarrow{\text{MgCl}}_{2}\text{(}aq\text{)}+\text{Ni(}s\text{)}
    2. PCl3(l)+Cl2(g)PCl5(s){\text{PCl}}_{3}\text{(}l\text{)}+{\text{Cl}}_{2}\text{(}g\text{)}\rightarrow{\text{PCl}}_{5}\text{(}s\text{)}
    3. C2H4(g)+3O2(g)2CO2(g)+2H2O(g){\text{C}}_{2}{\text{H}}_{4}\text{(}g\text{)}+3{\text{O}}_{2}\text{(}g\text{)}\rightarrow 2{\text{CO}}_{2}\text{(}g\text{)}+2{\text{H}}_{2}\text{O(}g\text{)}
    4. Zn(s)+H2SO4(aq)ZnSO4(aq)+H2(g)\text{Zn(}s\text{)}+{\text{H}}_{2}{\text{SO}}_{4}\text{(}aq\text{)}\rightarrow{\text{ZnSO}}_{4}\text{(}aq\text{)}+{\text{H}}_{2}\text{(}g\text{)}
    5. 2K2S2O3(s)+I2(s)K2S4O6(s)+2KI(s)2{\text{K}}_{2}{\text{S}}_{2}{\text{O}}_{3}\text{(}s\text{)}+{\text{I}}_{2}\text{(}s\text{)}\rightarrow{\text{K}}_{2}{\text{S}}_{4}{\text{O}}_{6}\text{(}s\text{)}+2\text{KI(}s\text{)}
    6. 3Cu(s)+8HNO3(aq)3Cu(NO3)2(aq)+2NO(g)+4H2O(l)3\text{Cu(}s\text{)}+8{\text{HNO}}_{3}\text{(}aq\text{)}\rightarrow 3\text{Cu}{\text{(}{\text{NO}}_{3}\text{)}}_{2}\text{(}aq\text{)}+2\text{NO(}g\text{)}+4{\text{H}}_{2}\text{O(}l\text{)}

  10. Complete and balance the following acid-base equations:

    1. HCl gas reacts with solid Ca(OH)2(s).
    2. A solution of Sr(OH)2 is added to a solution of HNO3.

  11. Complete and balance the following acid-base equations:

    1. A solution of HClO4 is added to a solution of LiOH.
    2. Aqueous H2SO4 reacts with NaOH.
    3. Ba(OH)2 reacts with HF gas.

  12. Complete and balance the following oxidation-reduction reactions, which give the highest possible oxidation state for the oxidized atoms.

    1. Al(s)+F2(g)\text{Al(}s\text{)}+{\text{F}}_{2}\text{(}g\text{)}\rightarrow
    2. Al(s)+CuBr2(aq)\text{Al(}s\text{)}+{\text{CuBr}}_{2}\text{(}aq\text{)}\rightarrow
      (single displacement)
    3. P4(s)+O2(g){\text{P}}_{4}\text{(}s\text{)}+{\text{O}}_{2}\text{(}g\text{)}\rightarrow
    4. Ca(s)+H2O(l)\text{Ca(}s\text{)}+{\text{H}}_{2}\text{O(}l\text{)}\rightarrow
      (products are a strong base and a diatomic gas)

  13. Complete and balance the following oxidation-reduction reactions, which give the highest possible oxidation state for the oxidized atoms.

    1. K(s)+H2O(l)\text{K(}s\text{)}+{\text{H}}_{2}\text{O(}l\text{)}\rightarrow
    2. Ba(s)+HBr(aq)\text{Ba(}s\text{)}+\text{HBr(}aq\text{)}\rightarrow
    3. Sn(s)+I2(s)\text{Sn(}s\text{)}+{\text{I}}_{2}\text{(}s\text{)}\rightarrow

  14. Complete and balance the equations for the following acid-base neutralization reactions. If water is used as a solvent, write the reactants and products as aqueous ions. In some cases, there may be more than one correct answer, depending on the amounts of reactants used.

    1. Mg(OH)2(s)+HClO4(aq)\text{Mg}{\text{(OH)}}_{2}\text{(}s\text{)}+{\text{HClO}}_{4}\text{(}aq\text{)}\rightarrow
    2. SO3(g)+H2O(l)(assume an excess of water and that the product dissolves){\text{SO}}_{3}\text{(}g\text{)}+{\text{H}}_{2}\text{O(}l\text{)}\rightarrow\phantom{\rule{0.3em}{0ex}}\text{(assume an excess of water and that the product dissolves)}
    3. SrO(s)+H2SO4(l)\text{SrO(}s\text{)}+{\text{H}}_{2}{\text{SO}}_{4}\text{(}l\text{)}\rightarrow

  15. When heated to 700–800 °C, diamonds, which are pure carbon, are oxidized by atmospheric oxygen. (They burn!) Write the balanced equation for this reaction.
  16. The military has experimented with lasers that produce very intense light when fluorine combines explosively with hydrogen. What is the balanced equation for this reaction?
  17. Write the molecular, total ionic, and net ionic equations for the following reactions:

    1. Ca(OH)2(aq)+HC2H3O2(aq)\text{Ca}{\text{(OH)}}_{2}\text{(}aq\text{)}+{\text{HC}}_{2}{\text{H}}_{3}{\text{O}}_{2}\text{(}aq\text{)}\rightarrow
    2. H3PO4(aq)+CaCl2(aq){\text{H}}_{3}{\text{PO}}_{4}\text{(}aq\text{)}+{\text{CaCl}}_{2}\text{(}aq\text{)}\rightarrow

  18. Great Lakes Chemical Company produces bromine, Br2, from bromide salts such as NaBr, in Arkansas brine by treating the brine with chlorine gas. Write a balanced equation for the reaction of NaBr with Cl2.
  19. In a common experiment in the general chemistry laboratory, magnesium metal is heated in air to produce MgO. MgO is a white solid, but in these experiments it often looks gray, due to small amounts of Mg3N2, a compound formed as some of the magnesium reacts with nitrogen. Write a balanced equation for each reaction.
  20. Lithium hydroxide may be used to absorb carbon dioxide in enclosed environments, such as manned spacecraft and submarines. Write an equation for the reaction that involves 2 mol of LiOH per 1 mol of CO2. (Hint: Water is one of the products.)
  21. Calcium propionate is sometimes added to bread to retard spoilage. This compound can be prepared by the reaction of calcium carbonate, CaCO3, with propionic acid, C2H5CO2H, which has properties similar to those of acetic acid. Write the balanced equation for the formation of calcium propionate.
  22. Complete and balance the equations of the following reactions, each of which could be used to remove hydrogen sulfide from natural gas:

    1. Ca(OH)2(s)+H2S(g)\text{Ca}{\text{(}\text{OH}\text{)}}_{2}\text{(}s\text{)}+{\text{H}}_{2}\text{S(}g\text{)}\rightarrow
    2. Na2CO3(aq)+H2S(g){\text{Na}}_{2}{\text{CO}}_{3}\text{(}aq\text{)}+{\text{H}}_{2}\text{S(}g\text{)}\rightarrow

  23. Copper(II) sulfide is oxidized by molecular oxygen to produce gaseous sulfur trioxide and solid copper(II) oxide. The gaseous product then reacts with liquid water to produce liquid hydrogen sulfate as the only product. Write the two equations which represent these reactions.
  24. Write balanced chemical equations for the reactions used to prepare each of the following compounds from the given starting material(s). In some cases, additional reactants may be required.

    1. solid ammonium nitrate from gaseous molecular nitrogen via a two-step process (first reduce the nitrogen to ammonia, then neutralize the ammonia with an appropriate acid)
    2. gaseous hydrogen bromide from liquid molecular bromine via a one-step redox reaction
    3. gaseous H2S from solid Zn and S via a two-step process (first a redox reaction between the starting materials, then reaction of the product with a strong acid)

  25. Calcium cyclamate Ca(C6H11NHSO3)2 is an artificial sweetener used in many countries around the world but is banned in the United States. It can be purified industrially by converting it to the barium salt through reaction of the acid C6H11NHSO3H with barium carbonate, treatment with sulfuric acid (barium sulfate is very insoluble), and then neutralization with calcium hydroxide. Write the balanced equations for these reactions.
  26. Complete and balance each of the following half-reactions (steps 2–5 in half-reaction method):

    1. Sn4+(aq)Sn2+(aq){\text{Sn}}^{\text{4+}}\text{(}aq\text{)}\rightarrow{\text{Sn}}^{\text{2+}}\text{(}aq\text{)}
    2. [Ag(NH3)2]+(aq)Ag(s)+NH3(aq){\left[\text{Ag}{\text{(}{\text{NH}}_{3}\text{)}}_{2}\right]}^{+}\text{(}aq\text{)}\rightarrow\text{Ag}\text{(}s\text{)}+{\text{NH}}_{3}\text{(}aq\text{)}
    3. Hg2Cl2(s)Hg(l)+Cl(aq){\text{Hg}}_{2}{\text{Cl}}_{2}\text{(}s\text{)}\rightarrow\text{Hg}\text{(}l\text{)}+{\text{Cl}}^{-}\text{(}aq\text{)}
    4. H2O(l)O2(in acidic solution){\text{H}}_{2}\text{O(}l\text{)}\rightarrow{\text{O}}_{2}\text{(}\text{in acidic solution}\text{)}
    5. IO3(aq)I2(s){\text{IO}}_{3}{}^{-}\text{(}aq\text{)}\rightarrow{\text{I}}_{2}\text{(}s\text{)}
    6. SO32-(aq)SO42-(aq) (in acidic solution){\text{SO}}_{3}{}^{\text{2-}}\text{(}aq\text{)}\rightarrow{\text{SO}}_{4}{}^{\text{2-}}\text{(}aq\text{) (in acidic solution)}
    7. MnO4(aq)Mn2+(aq) (in acidic solution){\text{MnO}}_{4}{}^{-}\text{(}aq\text{)}\rightarrow{\text{Mn}}^{\text{2+}}\text{(}aq\text{) (in acidic solution)}
    8. Cl(aq)ClO3(aq) (in basic solution){\text{Cl}}^{-}\text{(}aq\text{)}\rightarrow{\text{ClO}}_{3}{}^{-}\text{(}aq\text{) (in basic solution)}

  27. Complete and balance each of the following half-reactions (steps 2–5 in half-reaction method):

    1. Cr2+(aq)Cr3+(aq){\text{Cr}}^{\text{2+}}\text{(}aq\text{)}\rightarrow{\text{Cr}}^{\text{3+}}\text{(}aq\text{)}
    2. Hg(l)+Br(aq)HgBr42-(aq)\text{Hg(}l\text{)}+{\text{Br}}^{-}\text{(}aq\text{)}\rightarrow{\text{HgBr}}_{4}{}^{\text{2-}}\text{(}aq\text{)}
    3. ZnS(s)Zn(s)+S2-(aq)\text{ZnS(}s\text{)}\rightarrow\text{Zn(}s\text{)}+{\text{S}}^{\text{2-}}\text{(}aq\text{)}
    4. H2(g)H2O(l)(in basic solution){\text{H}}_{2}\text{(}g\text{)}\rightarrow{\text{H}}_{2}\text{O(}l\text{)}\text{(}\text{in basic solution}\text{)}
    5. H2(g)H3O+(aq)(in acidic solution){\text{H}}_{2}\text{(}g\text{)}\rightarrow{\text{H}}_{3}{\text{O}}^{\text{+}}\text{(}aq\text{)}\text{(}\text{in acidic solution}\text{)}
    6. NO3(aq)HNO2(aq)(in acidic solution){\text{NO}}_{3}{}^{-}\text{(}aq\text{)}\rightarrow{\text{HNO}}_{2}\text{(}aq\text{)}\text{(}\text{in acidic solution}\text{)}
    7. MnO2(s)MnO4(aq)(in basic solution){\text{MnO}}_{2}\text{(}s\text{)}\rightarrow{\text{MnO}}_{4}{}^{-}\text{(}aq\text{)}\text{(}\text{in basic solution}\text{)}
    8. Cl(aq)ClO3(aq) (in acidic solution){\text{Cl}}^{-}\text{(}aq\text{)}\rightarrow{\text{ClO}}_{3}{}^{-}\text{(}aq\text{) (in acidic solution)}

  28. Balance each of the following equations according to the half-reaction method:

    1. Sn2+(aq)+Cu2+(aq)Sn4+(aq)+Cu+(aq){\text{Sn}}^{\text{2+}}\text{(}aq\text{)}+{\text{Cu}}^{\text{2+}}\text{(}aq\text{)}\rightarrow{\text{Sn}}^{\text{4+}}\text{(}aq\text{)}+{\text{Cu}}^{+}\text{(}aq\text{)}
    2. H2S(g)+Hg22+(aq)Hg(l)+S(s) (in acid){\text{H}}_{2}\text{S(}g\text{)}+{\text{Hg}}_{2}{}^{\text{2+}}\text{(}aq\text{)}\rightarrow\text{H}g\text{(}l\text{)}+\text{S(}s\text{) (in acid)}
    3. CN(aq)+ClO2(aq)CNO(aq)+Cl(aq) (in acid){\text{CN}}^{-}\text{(}aq\text{)}+{\text{ClO}}_{2}\text{(}aq\text{)}\rightarrow{\text{CNO}}^{-}\text{(}aq\text{)}+{\text{Cl}}^{-}\text{(}aq\text{) (in acid)}
    4. Fe2+(aq)+Ce4+(aq)Fe3+(aq)+Ce3+(aq){\text{Fe}}^{\text{2+}}\text{(}aq\text{)}+{\text{Ce}}^{\text{4+}}\text{(}aq\text{)}\rightarrow{\text{Fe}}^{\text{3+}}\text{(}aq\text{)}+{\text{Ce}}^{\text{3+}}\text{(}aq\text{)}
    5. HBrO(aq)Br(aq)+O2(g) (in acid)\text{HBrO(}aq\text{)}\rightarrow{\text{Br}}^{-}\text{(}aq\text{)}+{\text{O}}_{2}\text{(}g\text{) (in acid)}

  29. Balance each of the following equations according to the half-reaction method:

    1. Zn(s)+NO3(aq)Zn2+(aq)+N2(g) (in acid)\text{Zn(}s\text{)}+{\text{NO}}_{3}{}^{-}\text{(}aq\text{)}\rightarrow{\text{Zn}}^{\text{2+}}\text{(}aq\text{)}+{\text{N}}_{2}\text{(}g\text{) (in acid)}
    2. Zn(s)+NO3(aq)Zn2+(aq)+NH3(aq) (in base)\text{Zn(}s\text{)}+{\text{NO}}_{3}{}^{-}\text{(}aq\text{)}\rightarrow{\text{Zn}}^{\text{2+}}\text{(}aq\text{)}+{\text{NH}}_{3}\text{(}aq\text{) (in base)}
    3. CuS(s)+NO3(aq)Cu2+(aq)+S(s)+NO(g) (in acid)\text{CuS(}s\text{)}+{\text{NO}}_{3}{}^{-}\text{(}aq\text{)}\rightarrow{\text{Cu}}^{\text{2+}}\text{(}aq\text{)}+\text{S(}s\text{)}+\text{NO(}g\text{) (in acid)}
    4. NH3(aq)+O2(g)NO2(g) (gas phase){\text{NH}}_{3}\text{(}aq\text{)}+{\text{O}}_{2}\text{(}g\text{)}\rightarrow{\text{NO}}_{2}\text{(}g\text{) (gas phase)}
    5. Cl2(g)+OH(aq)Cl(aq)+ClO3(aq) (in base){\text{Cl}}_{2}\text{(}g\text{)}+{\text{OH}}^{-}\text{(}aq\text{)}\rightarrow{\text{Cl}}^{-}\text{(}aq\text{)}+{\text{ClO}}_{3}{}^{-}\text{(}aq\text{) (in base)}
    6. H2O2(aq)+MnO4(aq)Mn2+(aq)+O2(g) (in acid){\text{H}}_{2}{\text{O}}_{2}\text{(}aq\text{)}+{\text{MnO}}_{4}{}^{-}\text{(}aq\text{)}\rightarrow{\text{Mn}}^{\text{2+}}\text{(}aq\text{)}+{\text{O}}_{2}\text{(}g\text{) (in acid)}
    7. NO2(g)NO3(aq)+NO2(aq) (in base){\text{NO}}_{2}\text{(}g\text{)}\rightarrow{\text{NO}}_{3}{}^{-}\text{(}aq\text{)}+{\text{NO}}_{2}{}^{-}\text{(}aq\text{) (in base)}
    8. Fe3+(aq)+I(aq)Fe2+(aq)+I2(aq){\text{Fe}}^{\text{3+}}\text{(}aq\text{)}+{\text{I}}^{-}\text{(}aq\text{)}\rightarrow{\text{Fe}}^{\text{2+}}\text{(}aq\text{)}+{\text{I}}_{2}\text{(}aq\text{)}

  30. Balance each of the following equations according to the half-reaction method:

    1. MnO4(aq)+NO2(aq)MnO2(s)+NO3(aq) (in base){\text{MnO}}_{4}{}^{-}\text{(}aq\text{)}+{\text{NO}}_{2}{}^{-}\text{(}aq\text{)}\rightarrow{\text{MnO}}_{2}\text{(}s\text{)}+{\text{NO}}_{3}{}^{-}\text{(}aq\text{) (in base)}
    2. MnO42-(aq)MnO4(aq)+MnO2(s) (in base){\text{MnO}}_{4}{}^{\text{2-}}\text{(}aq\text{)}\rightarrow{\text{MnO}}_{4}{}^{-}\text{(}aq\text{)}+{\text{MnO}}_{2}\text{(}s\text{) (in base)}
    3. Br2(l)+SO2(g)Br(aq)+SO42-(aq) (in acid){\text{Br}}_{2}\text{(}l\text{)}+{\text{SO}}_{2}\text{(}g\text{)}\rightarrow{\text{Br}}^{-}\text{(}aq\text{)}+{\text{SO}}_{4}{}^{\text{2-}}\text{(}aq\text{) (in acid)}


acid: substance that produces H3O+ when dissolved in water

acid-base reaction: reaction involving the transfer of a hydrogen ion between reactant species

base: substance that produces OH when dissolved in water

combustion reaction: vigorous redox reaction producing significant amounts of energy in the form of heat and, sometimes, light

half-reaction: an equation that shows whether each reactant loses or gains electrons in a reaction.

insoluble: of relatively low solubility; dissolving only to a slight extent

neutralization reaction: reaction between an acid and a base to produce salt and water

oxidation: process in which an element’s oxidation number is increased by loss of electrons

oxidation-reduction reaction: (also, redox reaction) reaction involving a change in oxidation number for one or more reactant elements

oxidation number: (also, oxidation state) the charge each atom of an element would have in a compound if the compound were ionic

oxidizing agent: (also, oxidant) substance that brings about the oxidation of another substance, and in the process becomes reduced

precipitate: insoluble product that forms from reaction of soluble reactants

precipitation reaction: reaction that produces one or more insoluble products; when reactants are ionic compounds, sometimes called double-displacement or metathesis

reduction: process in which an element’s oxidation number is decreased by gain of electrons

reducing agent: (also, reductant) substance that brings about the reduction of another substance, and in the process becomes oxidized

salt: ionic compound that can be formed by the reaction of an acid with a base that contains a cation and an anion other than hydroxide or oxide

single-displacement reaction: (also, replacement) redox reaction involving the oxidation of an elemental substance by an ionic species

soluble: of relatively high solubility; dissolving to a relatively large extent

solubility: the extent to which a substance may be dissolved in water, or any solvent

strong acid: acid that reacts completely when dissolved in water to yield hydronium ions

strong base: base that reacts completely when dissolved in water to yield hydroxide ions

weak acid: acid that reacts only to a slight extent when dissolved in water to yield hydronium ions

weak base: base that reacts only to a slight extent when dissolved in water to yield hydroxide ions

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