General Chemistry/Types of chemical reactions

The general form of a synthesis reaction is A + B → AB. Synthesis reactions "put things together".

Helpful Hint!
Synthesis reactions always yield one product. Reversing a synthesis reaction will give you a decomposition reaction.

$2{\hbox{H}}_{2(g)}+{\hbox{O}}_{2(g)}\to 2{\hbox{H}}_{2}{\hbox{O}}_{(l)}$	This is the most well-known example of a synthesis reaction—the formation of water via the combustion of hydrogen gas and oxygen gas.
$2{\hbox{Na}}_{(s)}+{\hbox{Cl}}_{2(g)}\to 2{\hbox{NaCl}}_{(s)}$	Another example of a synthesis reaction is the formation of sodium chloride (table salt).

Because of the very high reactivities of sodium metal and chlorine gas, this reaction releases a tremendous amount of heat and light energy. Recall that atoms release energy as they become stable, and consider the octet rule when determining why this reaction has such favorable features.

Decomposition Reactions

These are the same number of synthesis reactions, with the format AB → A + B. Decomposition reactions "take things apart". Just as synthesis reactions can only form one product, decomposition reactions can only start with one reactant. Compounds that are unstable decompose quickly without outside assistance. Also they help other atoms to decompose better.

$2{\hbox{H}}_{2}{\hbox{O}}_{(l)}{\xrightarrow {electricity}}2{\hbox{H}}_{2(g)}+{\hbox{O}}_{2(g)}$	One example is the electrolysis of water (passing water through electrical current) to form hydrogen gas and oxygen gas.
$2{\hbox{H}}_{2}{\hbox{O}}_{2(l)}\to 2{\hbox{H}}_{2}{\hbox{O}}_{(l)}+{\hbox{O}}_{2(g)}$	Hydrogen peroxide slowly decomposes into water and oxygen because it is somewhat unstable. The process is sped up by the energy from light, so hydrogen peroxide is often stored in dark containers to slow down the decomposition.
${\hbox{H}}_{2}{\hbox{CO}}_{3(aq)}\to {\hbox{H}}_{2}{\hbox{O}}_{(l)}+{\hbox{C}}{\hbox{O}}_{2(g)}$	Carbonic acid is the carbonation that is dissolved in soda. It decomposes into carbon dioxide and water, which is why an opened drink loses its fizz.

Single Displacement Reactions

Single displacement reaction, also called single replacement, is a reaction in which 2 elements are substituted for another element in a compound. The starting materials are always pure elements, such as a pure zinc metal or hydrogen gas plus an aqueous compound. When a displacement reaction occurs, a new aqueous compound and a different pure element are generated as products. Its format is AB + C → AC + B. Single Diplacement Adding hydrochloric acid to zinc will cause a gas to bubble out:

${\hbox{Zn}}_{(s)}+2{\hbox{HCl}}_{(aq)}\to {\hbox{ZnCl}}_{2(aq)}+{\hbox{H}}_{2(g)}$

Double Displacement Reactions

In these reactions, two compounds swap components, in the format AB + CD → AD + CB

This is also called an "exchange". Here are the examples below:

1.) HCl + NaOH ----> NaCl + H2O

Precipitation

A precipitation reaction occurs when an ionic substance comes out of solution and forms an insoluble (or slightly soluble) solid. The solid which comes out of solution is called a precipitate. This can occur when two soluble salts (ionic compounds) are mixed and form an insoluble one—the precipitate.

${\hbox{2Pb}}({\hbox{NO}}_{3})_{2(aq)}{\xrightarrow[{}]{\Delta }}{\hbox{2PbO}}_{(s)}+4{\hbox{NO}}_{2(aq)}+{O}_{2(g)}$	An example is lead nitrate mixed with potassium iodide, which forms a bright yellow precipitate of lead iodide.
${\hbox{Pb}}_{(aq)}^{2+}+2{\hbox{NO}}_{3(aq)}^{-}+2{\hbox{K}}_{(aq)}^{+}+2{\hbox{I}}_{(aq)}^{-}\to {\hbox{PbI}}_{2(s)}+2{\hbox{K}}_{(aq)}^{+}+2{\hbox{NO}}_{3(aq)}^{-}$	Note that the lead iodide is formed as a solid. The previous equation is written in molecular form, which is not the best way of describing the reaction. Each of the elements really exist in solution as individual ions, not bonded to each other (as in potassium iodide crystals). If we write the above as an ionic equation, we get a much better idea of what is actually happening.
${\hbox{Pb}}_{(aq)}^{2+}+2{\hbox{I}}_{(aq)}^{-}\to {\hbox{PbI}}_{2(s)}$	Notice the like terms on both sides of the equation. These are called spectator ions because they do not participate in the reaction. They can be ignored, and the net ionic equation is written.

In the solution, there exists both lead and iodide ions. Because lead iodide is insoluble, they spontaneously crystallise and form the precipitate.

Acid-Base Neutralization

In simple terms, an acid is a substance which can lose a H⁺ ion (i.e. a proton) and a base is a substance which can accept a proton. When equal amounts of an acid and base react, they neutralize each other, forming species which aren't as acidic or basic.

${\hbox{HCl}}_{(aq)}+{\hbox{NaOH}}_{(aq)}\to {\hbox{H}}_{2}{\hbox{O}}_{(l)}+{\hbox{NaCl}}_{(aq)}$	For example, when hydrochloric acid and sodium hydroxide react, they form water and sodium chloride (table salt).
${\hbox{H}}_{(aq)}^{+}+{\hbox{OH}}_{(aq)}^{-}\to {\hbox{H}}_{2}{\hbox{O}}_{(l)}$	Again, we get a clearer picture of what's happening if we write a net ionic equation.

Acid base reactions often happen in aqueous solution, but they can also occur in the gaseous state. Acids and bases will be discussed in much greater detail in the acids and bases section.

Combustion

The combustion of methane (releasing heat and light)

Combustion, better known as burning, is the combination of a substance with oxygen. The products are carbon dioxide, water, and possible other waste products. Combustion reactions release large amounts of heat. C₃H₈, better known as propane, undergoes combustion. The balanced equation is:

${\hbox{C}}_{3}{\hbox{H}}_{8}+5{\hbox{O}}_{2}\to 3{\hbox{CO}}_{2}+4{\hbox{H}}_{2}{\hbox{O}}$

Combustion is similar to a decomposition reaction, except that oxygen and heat are required for it to occur. If there is not enough oxygen, the reaction may not occur. Sometimes, with limited oxygen, the reaction will occur, but it produces carbon monoxide (CO) or even soot. In that case, it is called incomplete combustion. If the substances being burned contain atoms other than hydrogen and oxygen, then waste products will also form. Coal is burned for heating and energy purposes, and it contains sulfur. As a result, sulfur dioxide is released, which is a pollutant. Coal with lower sulfur content is more desirable, but more expensive, because it will release less of the sulfur-based pollutants.

Organic Reactions

This is carboxylic acid. All functional groups end with an "R"—a placeholder for the rest of the molecule.

Organic reactions occur between organic molecules (molecules containing carbon and hydrogen). Since there is a virtually unlimited number of organic molecules, the scope of organic reactions is very large. However, many of the characteristics of organic molecules are determined by functional groups—small groups of atoms that react in predictable ways.

Another key concept in organic reactions is Lewis basicity. Parts of organic molecules can be electrophillic (electron-loving) or nucleophillic (nucleus, or positive loving). Nucleophillic regions have an excess of electrons—they act as Lewis bases—whereas electrophillic areas are electron deficient and act as Lewis acids. The nucleophillic and electrophillic regions attract and react with each other. Organic reactions are beyond the scope of this book, and are covered in more detail in Organic Chemistry. However, most organic substances can undergo replacement reactions and combustion reactions, as you have already learned.

Redox

The formation of hydrogen fluoride from the elements requires reduction of fluorine and oxidation of hydrogen.

Redox is an abbreviation of reduction/oxidation reactions. This is exactly what happens in a redox reaction, one species is reduced and another is oxidized. Reduction involves a gain of electrons and oxidation involves a loss, so a redox reaction is one in which electrons are transferred between species. Reactions where something is "burnt" (burning means being oxidised) are examples of redox reactions, however, oxidation reactions also occur in solution, which is very useful and forms the basis of electrochemistry.

Redox reactions are often written as two half-reactions showing the reduction and oxidation processes separately. These half-reactions are balanced (by multiplying each by a coefficient) and added together to form the full equation. When magnesium is burnt in oxygen, it loses electrons (it is oxidised). Conversely, the oxygen gains electrons from the magnesium (it is reduced).

${\begin{matrix}{\hbox{Mg}}&\to &{\hbox{Mg}}^{2+}+2e^{-}&\times 2\\{\hbox{O}}_{2}+4e^{-}&\to &2{\hbox{O}}^{2-}&\times 1\\2{\hbox{Mg}}+{\hbox{O}}_{2}+4e^{-}&\to &2{\hbox{MgO}}+4e^{-}&\ \\\end{matrix}}$

Redox reactions will be discussed in greater detail in the redox section.