Search 75,858 tutors
FIND TUTORS

Types of Chemical Reactions

Single Displacement Reactions

Form: AB + C --> A + CB

Example: Zn(OH)2 + 2 Na --> Zn (s) + 2 NaOH

Notice that in this reaction, the hydroxide group moves from the Zinc (Zn) to the Sodium (Na). We call this reaction a single displacement because only one ion is moving.

Double Displacement Reactions

Form: AB + CD --> AD + BC

Example: AgNO3 + HCl --> AgCl (s) + HNO3

Notice that in this reaction, the nitrate ion (NO3) and the chloride ion (Cl) trade places within the equation. The fact that two ions are moving causes it to be named a double displacement reaction.

Precipitation Reactions

Form: can be single or double displacement reactions (see above) that produce a solid
AB + C --> A (s) + CB or AB + C --> A + CB (s)
or, AB + CD --> AD + BC (s) or AB + CD --> AD (s) + BC

Example: 3 NaOH + Al --> 3 Na + Al(OH)3 (s) (single displacement precipitation)
NaCl + AgNO3 --> NaNO3 + AgCl (s) (double displacement precipitation)

Notice that in these types of reactions, we denote a certain element or compound with (s), which stands for "solid". This means that one of the products is not soluble; it cannot dissolve in solution. When an element or a compound is not soluble, the element or compound precipitates, meaning it becomes solid and sinks to the bottom of the solution. Therefore, it is no longer a part of the solution; it exists, but it is not mixed thoroughly with the solution. Rather, it is a solid sitting within the solution. Many reactions have at least one element or compound which is aqueous (aq), meaning that it dissolves in water. However, some solutions also have an element or compound that is not aqueous (therefore not soluble) in water, which is what produces a precipitate.

There are rules for telling whether or not an element or compound will precipitate: these are called Solubility Rules. They are as follows:

1. Salts containing Group I elements are soluble (Li+, Na+, K+, Cs+, Rb+). Exceptions to this rule are rare. Salts containing the ammonium ion (NH4+) are also soluble.
2. Salts containing nitrate ion (NO3-) are generally soluble.
3. Salts containing Cl-, Br-, I- are generally soluble. Important exceptions to this rule are halide salts of Ag+, Pb2+, and (Hg2)2+. Thus, AgCl, PbBr2, and Hg2Cl2 are all insoluble.
4. Most silver salts are insoluble. AgNO3 and Ag(C2H3O2) are common soluble salts of silver; virtually anything else is insoluble.
5. Most sulfate salts are soluble. Important exceptions to this rule include BaSO4, PbSO4, Ag2SO4 and SrSO4.
6. Most hydroxide salts are only slightly soluble. Hydroxide salts of Group I elements are soluble. Hydroxide salts of Group II elements (Ca, Sr, and Ba) are slightly soluble. Hydroxide salts of transition metals and Al3+ are insoluble. Thus, Fe(OH)3, Al(OH)3, Co(OH)2 are not soluble.
7. Most sulfides of transition metals are highly insoluble. Thus, CdS, FeS, ZnS, Ag2S are all insoluble. Arsenic, antimony, bismuth, and lead sulfides are also insoluble.
8. Carbonates are frequently insoluble. Group II carbonates (Ca, Sr, and Ba) are insoluble. Some other insoluble carbonates include FeCO3 and PbCO3.
9. Chromates are frequently insoluble. Examples: PbCrO4, BaCrO4
10. Phosphates are frequently insoluble. Examples: Ca3(PO4)2, Ag3PO4
11. Fluorides are frequently insoluble. Examples: BaF2, MgF2, PbF2.

These are known as the solubility rules. Although you may find a slightly different set of rules if you were to look them up in a text or other resource, they are generally accepted as listed above, with few, rare variations. You should keep these rules with you whenever working with chemical equations in order to discern which solutions will have precipitates and which solutions will not.

Combustion

Form: CxHy + O2 --> CO2 + H2O + energy

Example: CH4 + 2O2 --> CO2 + 2H2O + energy

Combustion is also known as burning. It always includes O2 (g) as a reactant-just like fire needs oxygen to burn, combustion reactions need oxygen to occur. Combustion reactions also produce heat (released as energy). You can learn more about this in the Thermodynamics (link) help section. Combustion reactions usually have CxHy as a reactant and have CO2 and H2O as products. Combustion reactions almost never go to completion. Practice problems usually assume that combustion does go into completion; however, when these reactions occur in nature, they reach equilibrium before they reach completion.

The following reaction is another example of combustion:

2 H2 + O2 --> 2H2O (g) + heat

This reaction is an example of a combustion reaction that does not include carbon. In this example, the product is gaseous water, or water vapor (also known commonly as steam).

Acid/Base Reactions

Form: acid+ + base- --> salt + water

Example: H2SO4 + 2 NaOH --> Na2SO4 + 2H2O

More generally, acid base reactions H+ + OH- --> H2O, or (now, more commonly): H3O+ +OH- --> H2O

Acid base reactions are neutralization reactions. These reactions occur when an acid (commonly with a positive hydrogen cation) and a base (commonly with a negative hydroxide anion) react and bond together to form water. Usually, the other components react to form a salt, like NaCl.

Bronsted-Lowry Acids and Bases

It should be noted that in the study of chemistry, there are differing definitions of acids and bases. However, the most commonly accepted is the Bronsted-Lowry model of acids and bases, which was composed in 1923 by Johannes Nicolaus Bronsted and Martin Lowry. These two scientists did not collaborate; rather, they worked independently of each other and came up with the same theory. Thus, the theory was accepted as true and right. Their model does not identify acids and bases solely based on the formulation of a salt and a solvent (water) as we described above. Instead, they define acids as "compounds having the ability to donate a proton" and bases as "compounds having the ability to receive a proton." The proton is never donated or received directly from the nucleus of an atom; rather, it is through the addition or removal of a hydrogen atom (H+) from a compound.

Here's a simple way of looking at conjugate acids and bases.

A conjugate acid is formed by the addition of a proton in the form of a H+ ion. A common example of a conjugate acid is ammonium, NH4+. The ammonium ion starts out as NH3, a compound consisting of hydrogen and nitrogen, and accepts another hydrogen ion, making NH4+. Now, the ammonium ion is considered to have an "extra" hydrogen ion that it is able to donate to a compound in need of a proton.

The reaction for this conjugate acid is: NH3 + H+ --> NH4+ (The product NH4+ would be the conjugate acid)

A conjugate base is formed by the removal of a proton in the form of a H+ ion. A common example of a conjugate base is the chloride ion, Cl-. The chloride ion would start as HCl (hydrochloric acid), a compound consisting of a hydrogen ion and a chloride ion bonded together. When this compound gives up the hydrogen ion (H+) it is simply left with the chloride ion, Cl-. Now, chloride is ready to accept another proton in the form of a hydrogen ion.

The reaction for this conjugate base is: HCl --> H+ + Cl- (The product Cl- would be the conjugate base)

Redox Reactions (Oxidation-Reduction Reactions)

Form: X --> X+ + e- (oxidation)
Y + e- --> Y- (reduction)

Oxidation-reduction reactions, known commonly as redox reactions, are reactions in which the oxidation state (oxidation number) of atoms changes. More often, electron transfers between atoms identify these types of reactions. Oxidation is the loss of electrons, which increases an atom's oxidation state. Reduction is the gain of electrons, which decreases an atom's oxidation state. On paper, these are fine definitions; however, in nature, the transfer of electrons may never actually occur. For example, in reactions with covalent bonds, it is very possible for the oxidation state to change while never actually experiencing a transfer in electrons. Therefore, it is better to refer to oxidation as simply an increase in oxidation state and reduction as a decrease in oxidation state. The following is an example of a redox reaction:

H2 + F2 --> 2 HF

Redox reactions can be split up into two smaller equations, like this:

H2 ? 2 H+ + 2 e-

and

F2 + 2 e- ? 2 F-

The first smaller equation shows the oxidation of hydrogen. The second smaller equation shows the reduction of fluorine. Notice that the first reaction splits elemental hydrogen into hydrogen ions and electrons, making the hydrogen atom more positive. Also notice that the second reaction combines elemental fluorine with two electrons, which reduce fluorine by making it more negative. Therefore, when we talk about this reaction, we can say the following:

Hydrogen is being oxidized.

Fluorine is being reduced.

Fluorine is the oxidizing agent, because it oxidizes hydrogen by accepting two electrons and is, in turn, reduced.

Hydrogen is the reducing agent, because it reduces fluorine by giving it two electrons and is, in turn, oxidized.

This may seem really backwards, and from a very basic perspective, it is. The naming of these does not make good sense, and it can be really tricky to call them the right things.

Very simply:

The element that is oxidized becomes more positive because it loses its electrons. The element that is oxidized is known as the reducing agent, because it is helping the other element be reduced.

The element that is reduced becomes more negative because it gains electrons. The element that is reduced is known as the oxidizing agent, because it is helping the other element be oxidized.

Let's think about this. In the above example, hydrogen is giving up two electrons. When you get rid of electrons, your element becomes more positive, so we can say that element is oxidized. However, hydrogen is giving the two electrons to fluorine, which is making fluorine more negative than when we started out. Therefore, fluorine is reduced (made more negative). Now, since hydrogen is helping to reduce fluorine, hydrogen is our reducing agent. The element that gives up electrons is always the reducing agent. Since fluorine is accepting the electrons, it is helping to oxidize hydrogen. Therefore, fluorine is the oxidizing agent.

Please make a mental note of this now. The reducing agent is oxidized and the oxidizing agent is reduced. They are absolutely switched around, and you have to remember this or your completion of redox reactions will seem very confusing.

Summary of Chemical Reactions

To sum it all up, there are many different types of reactions that occur in nature. They can be defined as follows:

1. Single Displacement reactions

Form: AB + C --> A + CB

2. Double Displacement reactions

Form: AB + CD --> AD + BC

3. Precipitation reactions

Form: can be single or double displacement reactions (see above) that produce a solid

4. Combustion

Form: CxHy + O2 --> CO2 + H2O + energy

5. Acid/base reactions

Form: acid+ + base- --> salt + water

6. Redox reactions (Oxidation-reduction reactions)

Form: X --> X+ + e- (oxidation)
Y + e- --> Y- (reduction)

Sign up for free to access more Science resources like . WyzAnt Resources features blogs, videos, lessons, and more about Science and over 250 other subjects. Stop struggling and start learning today with thousands of free resources!