A discussion of different classifications of chemical reactions in AP® Chemistry.

Chemical reactions occur when one or more molecules change chemically. In other words, bonds break between the reactants, and atoms rearrange to form new chemical bonds. These reconfigured atoms can either be a compound or just an element. The products have different chemical identities compared to the reactants. Common examples of chemical reactions include the burning of fuel, respiration, and the formation of rust. We can classify chemical reactions according to the nature of the reactants or how the products are formed. In AP® Chemistry, you will encounter problems that ask you to identify and predict the products of chemical reactions. In this post, we will discuss how to classify chemical reactions in detail.

Synthesis Reactions

In a synthesis reaction, two or more simple substances join to form more complicated products. This type of chemical reaction usually involves the formation of a binary compound. Binary compounds contain two different elements. The best way to recognize a synthesis reaction is to see if the reactants “combine” to form the product. For this reason, synthesis reactions are also called the direct combination reactions. The general form of a synthesis reaction is given as:

A+B\rightarrow AB

The most classic example of a synthesis reaction is the formation of water by reacting hydrogen and oxygen. Hydrogen gas burns in the presence of oxygen, which produces steam as a product, along with the release of heat. Hydrogen and oxygen are diatomic molecules, and the balanced equation for this reaction is given below. As we can see, there is only one product consisting of the two reactant elements.

2{ H }_{ 2 }\left( g \right) +{ O }_{ 2 }\left( g \right) \rightarrow 2{ H }_{ 2 }O\left( g \right)

Another type of synthesis reaction is the addition of metal to nonmetal elements. An example of this type is the formation of ferrous sulfide. Ferrous sulfide is used mainly in the stainless steel and alloy industries and is formed by heating a mixture of iron and sulfur powders. The reaction is exothermic.

8{ Fe }\left( s \right) +S_8\left( s \right) \rightarrow 8{ Fe }S\left( s \right)

An important example of a synthesis reaction is the photosynthesis process. Photosynthesis is the process wherein plants absorb carbon dioxide to produce oxygen and glucose. In this reaction, the reactants are not elements, as in the examples above; rather, they are compounds. Another difference is that this reaction generates two products. Although there are multiple products, we can see that the elements from the reactants combine to form glucose, one of the products.

6{ CO }_{ 2 }+6{ H }_{ 2 }O\rightarrow C_6{ H }_{ 12 }{ O }_{ 6 }+{ O }_{ 2 }

Other examples of synthesis reaction are given below.

1. Formation of aluminum oxide

4Al\left( s \right) +3{ O }_{ 2 }\left( g \right) \rightarrow 2{ Al }_{ 2 }{ O }_{ 3 }\left( s \right)

2. Synthesis of ammonia (Haber Process)

3{ H }_{ 2 }\left( g \right) +{ N }_{ 2 }\left( g \right) \rightarrow 2{ NH }_{ 3 }\left( g \right)

3. Formation of potassium chloride

2K\left( s \right) +{ Cl }_{ 2 }\left( g \right) \rightarrow 2KCl\left( s \right)

4. Formation of table salt

2Na\left( s \right) +C{ l }_{ 2 }\left( g \right) \rightarrow 2NaCl\left( s \right)

5. Synthesis of carbon dioxide

2CO\left( g \right) +{ O }_{ 2 }\left( g \right) \rightarrow 2C{ O }_{ 2 }\left( g \right)

Decomposition Reactions

A decomposition reaction typically involves the separation of a compound into simpler chemical species. The products are usually the constituent elements of the compounds. There are also decomposition reactions for which the products are simpler compounds compared to the reactants. Decomposition reactions are also called analysis reactions or breakdown reactions. In decomposition reactions, chemical bonds are broken. Compounds can easily undergo decomposition when exposed to some kind of external stress, such as humidity, heat, the presence of electricity, and change in acidity. This type of reaction is useful in analytical chemistry techniques, including gravimetric analysis and mass spectrometry. Decomposition reactions are the opposite of synthesis reactions, as evident in the general form for decomposition reactions:

AB\rightarrow A+B

One example of a decomposition reaction is the electrolysis of water. This reaction takes place when an electric current passes through water, which results in the dissociation of the water compound into diatomic molecules of oxygen and hydrogen. This is a common experiment in AP® Chemistry labs. The reaction for the electrolysis of water shows that for every mole of water, one mole of hydrogen gas and half a mole of oxygen gas is produced.

{ H }_{ 2 }O\left( g \right) \rightarrow { H }_{ 2 }\left( g \right) +\dfrac { 1 }{ 2 } { O }_{ 2 }\left( g \right)

Another industrial application of decomposition is the extraction of metals from their oxide ores. Metals are purified with the use of heat in a process called thermal decomposition. It is possible to decompose any metal provided enough heat input, which sometimes results in extremely high temperatures. The product for this process will always be the elemental form of the metal and oxygen gas. For example, mercuric oxide will decompose at a temperature above500^{ \circ }\text{C} through the following reaction:

2HgO\rightarrow 2Hg\left( l \right) +{ O }_{ 2 }\left( g \right)

Carbonate decomposition is also prominent in industries. When a carbonate metal decomposes, the products will be carbon dioxide and a metal oxide. Metals that are higher on the activity series will need more energy for their carbonates to decompose. Some carbonate decomposition reactions include the following:

1. { H }_{ 2 }{ CO }_{ 3 }\rightarrow { CO }_{ 2 }+{ H }_{ 2 }O
2. Cu{ CO }_{ 3 }\rightarrow { CO }_{ 2 }+CuO
3. { K }_{ 2 }{ CO }_{ 3 }\rightarrow { CO }_{ 2 }+{ K }_{ 2 }O
4. Ca{ CO }_{ 3 }\rightarrow { CO }_{ 2 }+CaO
5. Mg{ CO }_{ 3 }\rightarrow { CO }_{ 2 }+MgO

Combustion Reactions

We encounter combustion reactions often in our daily lives. This type of reaction involves a reaction between a fuel and an oxidizer. We will only consider organic combustions, where the fuel will always be a hydrocarbon and the oxidizer is oxygen. A hydrocarbon is a compound consisting mainly of carbon (C) and hydrogen (H). Combustion reactions are exothermic reactions, or reactions that release heat. There are two types of combustion: complete and incomplete combustions.

Complete combustion occurs when a sufficient amount of fuel is burned with enough oxygen. The products for complete combustion are always carbon dioxide and water. It is also known as “clean combustion” because it does not generate harmful compounds. On the other hand, incomplete combustion, also called “dirty combustion,” produces products other than carbon dioxide. This type of reaction usually occurs when there is a lack of oxygen, and the additional products contribute to pollution.

In AP® Chemistry, it is easy to recognize a combustion reaction. If the reactants are oxygen and a hydrocarbon, then the reaction is a combustion. Balancing a combustion reaction is where the difficulty lies. Balancing complete combustion reactions involves balancing the carbon and hydrogen first and then balancing the oxygen.

Now, we will consider the reaction of butane ({ C }_{ 4 }{ H }_{ 10 }), a common fuel for cigarette lighters and portable stoves:

{ C }_{ 4 }{ H }_{ 10 }+{ O }_{ 2 }\rightarrow { CO }_{ 2 }+{ H }_{ 2 }O

Since butane has four carbons, we need to put a coefficient of 4 for the { CO }_{ 2 } compound. Then, in order to balance the ten atoms of hydrogen on the reactant side, we need to place a coefficient of 5 with the water compound on the product side. Now, we need to balance the oxygen on both sides.

{ C }_{ 4 }{ H }_{ 10 }+{ O }_{ 2 }\rightarrow { 4CO }_{ 2 }+{ 5H }_{ 2 }O

On the product side, there is a total of 13 atoms of oxygen, such that the oxygen compound needs a coefficient of \dfrac { 13 }{ 2 }.

{ C }_{ 4 }{ H }_{ 10 }+\dfrac { 13 }{ 2 } { O }_{ 2 }\rightarrow { 4CO }_{ 2 }+{ 5H }_{ 2 }O

Since the coefficient of oxygen is a fraction, we need to multiply the whole equation by 2.

As a result, the balanced reaction is:

{ 2C }_{ 4 }{ H }_{ 10 }+13{ O }_{ 2 }\rightarrow { 8CO }_{ 2 }+{ 10H }_{ 2 }O

Below are some additional examples of balanced combustion reactions:

1. Combustion of propane: { C }_{ 3 }{ H }_{ 8 }+5{ O }_{ 2 }\rightarrow 3{ CO }_{ 2 }+4{ H }_{ 2 }O
2. Combustion of ethane: { 2C }_{ 2 }{ H }_{ 6 }+7{ O }_{ 2 }\rightarrow 4{ CO }_{ 2 }+6{ H }_{ 2 }O
3. Combustion of pentane: { C }_{ 5 }{ H }_{ 12 }+8{ O }_{ 2 }\rightarrow 5{ CO }_{ 2 }+6{ H }_{ 2 }O

Single Replacement Reactions

When one element substitutes another element in a compound, the reaction is a single replacement reaction (also called single displacement reaction). In this type of reaction, the reactants will always be a pure element and a compound. The general form of this type of reaction is:

A+BC\rightarrow AC+B

We can see that A replaces element B in compound BC, producing the pure element B as a product. Since C does not participate in the reaction, it is called the spectator ion. As the name suggests, C can be an element or an ion.

The activity series of metals limits the possibility of some reactions. The activity series is a guide for the reactivity of elements and helps you predict the products of replacement reactions. The table below shows the activity series of metals and halogens. The elements higher in the table are more reactive than the elements below them. More reactive elements can replace less reactive elements in a reaction. For example, calcium can replace magnesium, but not potassium, in a reaction. The rule also applies to halogens.

Table 1: Activity Series

Below, we will illustrate some examples of how to use the activity series to analyze reactions.

1. In the reaction of copper (Cu) and silver nitrate (Ag{ NO }_{ 3 }), nitrate ({ NO }_{ 3 }^{ - }) is the spectator ion. The metals we need to compare are silver and copper. Checking the activity series table, copper is higher compared to silver, which means that copper is more reactive than silver. Copper will replace silver, and the reaction is possible.

Cu\left( s \right) +Ag{ NO }_{ 3 }\left( aq \right) \rightarrow Ag\left( s \right) +Cu{ NO }_{ 3 }\left( aq \right)

2. Considering the reaction between zinc (Zn) and hydrochloric acid (HCl), the chlorine ion is the spectator ion. The table shows that zinc is more reactive than hydrogen. The reaction can occur, and zinc will replace hydrogen. Since hydrogen is a diatomic element, its stable form is { H }_{ 2 }\left( g \right), which must be one of the products for the reaction. The balanced reaction is given as:

Zn+2HCl\rightarrow Zn{ Cl }_{ 2 }+{ H }_{ 2 }

3. For the addition of bromine gas ({ Br }_{ 2 }) to sodium chloride (NaCl), the spectator ion is the cation, which is sodium. As we can see from the table above, bromine is located below chlorine, which means that it is less reactive than chlorine. This indicates that the reaction is not possible.

{ Br }_{ 2 }+NaCl\rightarrow \text{ No reaction possible}

4. With the reaction of chlorine gas and magnesium iodide, magnesium is the spectator ion. We need to compare chlorine and iodine on the halogen activity series. Chlorine is higher, and therefore, it can replace iodine. The balanced reaction is shown below.

{ Cl }_{ 2 }+{ Mgl }_{ 2 }\rightarrow { MgCl }_{ 2 }+{ l }_{ 2 }

5. Consider the reaction of iron and magnesium sulfate. Comparing iron and magnesium, iron is below magnesium, so it can’t replace magnesium in the reaction. This means that the reaction is not possible.

Fe+{ MgSO }_{ 4 }\rightarrow \text{ No reaction possible}

Double Replacement Reactions

The last classification for chemical reactions is the double replacement reaction. Put simply, the ions involved in this type of reaction change places to form two new compounds as products. This type of chemical reaction is also called metathesis. The general form for this reaction is:

AB+CD\rightarrow AD+BC

In the general form, { A }^{ + } and { C }^{ + } are the cations, and { B }^{ - } and { D }^{ - } are the anions. Switching the pairings of cations and anions will form two new compounds, AD and BC.

There are two types of double replacement reactions: precipitation reactions and neutralization reactions. Precipitation reactions involve two aqueous compounds that form a solid precipitate and a new aqueous compound as the products. Meanwhile, neutralization reactions concern reactions between acids and bases. If one of the reactants involved in a neutralization reaction is water, one of the products is a salt.

To know whether one of the products is a precipitate, we need to use the solubility guidelines. The solubility guidelines are a set of rules that state whether a pair of ions is soluble in water or not. If the compound is soluble, then it will not form a precipitate. The solubility rules are as follows:

1. Salts containing Group I elements or an ammonium cation are soluble.
2. Salts with nitrate anions are soluble.
3. Salts with chlorine, iodine, and bromine anions are generally soluble; however, if paired with cations { Ag }^{ + }, { Pb }_{ 2 }^{ + }, and { Hg }_{ 22 }^{ + }, they will precipitate.
4. Silver salts are generally insoluble except for silver nitrate and silver acetate.
5. Most sulfate salts are soluble except for sulfates of copper, barium, lead (II), silver, and strontium.
6. Hydroxide salts of Group I elements are soluble. Hydroxide salts of Group II elements are slightly soluble. Hydroxide salts formed with transition elements and aluminum (III) are insoluble.
7. Sulfides of transition metals are insoluble.
8. Carbonates, chromates, phosphates, and fluorides are insoluble.

Some examples of double replacement reactions are presented below:

1. In the addition of hydrochloric acid to silver nitrate, the cations are silver and potassium while the anions are nitrate and chlorine. Both sets of ions will interchange, forming two new compounds: silver chloride and potassium nitrate. Silver chloride is a white solid precipitate, which indicates that this reaction is a precipitation reaction. Silver chloride is used in photographic papers and stained glass manufacturing.

Ag{ NO }_{ 3 }\left( aq \right) +KCl\left( aq \right) \rightarrow AgCl\left( s \right) +{ KNO }_{ 3 }\left( aq \right)

2. A classic example of neutralization reaction is between a strong acid, hydrochloric acid, and a strong base, sodium hydroxide. In this reaction, they will form a neutral salt, sodium chloride.

HCl\left( aq \right) +NaOH\left( aq \right) \rightarrow NaCl\left( aq \right) +{ H }_{ 2 }O\left( l \right)

3. In the reaction between copper sulfate and sodium hydroxide, a blue precipitate is formed. This precipitate is the compound copper hydroxide. Different transition metals yield colored precipitates. This is apparent in the image above.

Cu{ SO }_{ 4 }\left( aq \right) +NaOH\left( aq \right) \rightarrow Cu{ \left( OH \right) }_{ 2 }\left( s \right) +{ Na }_{ 2 }{ SO }_{ 4 }\left( aq \right)

Other examples for double replacement reactions include the following:

1. Fe{ \left( OH \right) }_{ 3 }\left( aq \right) +{ NHO }_{ 3 }\left( aq \right) \rightarrow Fe{ NO }_{ 3 }\left( aq \right) +{ H }_{ 2 }O\left( l \right) (Neutralization Reaction)
2. 2Ag{ NO }_{ 3 }\left( aq \right) +{ Na }_{ 2 }S\left( aq \right) \rightarrow Ag2S\left( s \right) +2Na{ NO }_{ 3 }\left( aq \right) (Precipitation Reaction)
3. { HC }_{ 2 }{ H }_{ 3 }{ O }_{ 2 }\left( aq \right) +NaOH\left( aq \right) \rightarrow Na{ C }_{ 2 }{ H }_{ 3 }{ O }_{ 2 }\left( aq \right) +{ H }_{ 2 }O\left( l \right) (Neutralization Reaction)

When you study AP® Chemistry, you will always encounter a chemical reaction. We can use the table below as a simple guide on how to classify chemical reactions.

Table 2: Chemical Reactions Summary

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