Types and Applications of Chemical Reactions

For example, when sodium ($Na$) and chlorine ($Cl_2$) react, they form sodium chloride ($NaCl$), commonly known as table salt: $2Na + Cl_2 \longrightarrow 2NaCl$

2. Decomposition

Decomposition is the opposite of synthesis. A single compound breaks down into two or more simpler substances. Because breaking bonds requires energy, these reactions often need an external input such as heat, light, or electricity.

The general form is: $AB \longrightarrow A + B$

A practical example is the heating of baking soda ($NaHCO_3$). When heat is applied, it decomposes to release carbon dioxide ($CO_2$) and water ($H_2O$), which helps dough rise during baking: $2NaHCO_3 \xrightarrow{\Delta} Na_2CO_3 + CO_2 + H_2O$

3. Single Replacement (Displacement)

In a single replacement reaction, a more reactive element takes the place of another element within a compound. This is often used in industrial applications, such as galvanizing steel to prevent corrosion.

The general form is: $A + BC \longrightarrow AC + B$

For instance, when zinc ($Zn$) is placed in hydrochloric acid ($HCl$), the zinc replaces the hydrogen, resulting in zinc chloride ($ZnCl_2$) and hydrogen gas ($H_2$): $Zn + 2HCl \longrightarrow ZnCl_2 + H_2$

4. Double Replacement (Metathesis)

Double replacement reactions involve two ionic compounds exchanging ions to form two entirely new compounds. These reactions often result in the formation of a precipitate (an insoluble solid), a gas, or water.

The general form is: $AB + CD \longrightarrow AD + CB$

An everyday example is the reaction between soap and the calcium ions found in hard water, which produces insoluble "soap scum".

5. Combustion

Combustion is a rapid reaction between a fuel (typically a hydrocarbon) and oxygen gas ($O_2$). This reaction releases significant energy in the form of heat and light. The standard products of a hydrocarbon combustion are carbon dioxide ($CO_2$) and water ($H_2O$).

The general form is: $\text{Hydrocarbon} + O_2 \longrightarrow CO_2 + H_2O + \text{energy}$

A simple example is burning charcoal ($C$) in oxygen to produce carbon dioxide and heat: $C + O_2 \longrightarrow CO_2 + \text{heat}$

Representing Reactions: Chemical Equations

To communicate these transformations clearly, scientists use chemical equations.

Anatomy of an Equation

In a standard equation, the reactants are written on the left side, and the products are written on the right side, separated by a reaction arrow ($\longrightarrow$) that indicates the direction of the transformation.

To provide more detail, phase symbols are used to denote the physical state of each substance:

• $(s)$ for solid
• $(l)$ for liquid
• $(g)$ for gas
• $(aq)$ for aqueous (a substance dissolved in water)

Coefficients vs. Subscripts

It is crucial to distinguish between two types of numbers used in chemical formulas to avoid errors when balancing equations:

1. Subscripts: These are the small numbers located inside a chemical formula (e.g., the $_2$ in $H_2O$). They define the chemical identity and the specific ratio of atoms within a compound. Never change a subscript when balancing an equation, as doing so changes the substance itself.
2. Coefficients: These are the large numbers placed in front of a chemical formula (e.g., the $2$ in $2H_2O$). They indicate the relative number of molecules or units involved in the reaction. Coefficients are adjusted to balance the equation and satisfy the Law of Conservation of Mass.

Exam Tip: If you are asked to balance an equation, only touch the coefficients. If you change a subscript, you are no longer describing the same chemical reaction!