4 Chemical Reactions and Equations

Reactions and chemical equations

A chemical reaction changes substances called reactants into new substances called products. As chemical bonds break and form, atoms are rearranged; in an ordinary chemical reaction, atoms are not created or destroyed.

A chemical equation represents this change, with the reactants on the left and products on the right:

For example, $\mathrm{2H_2 + O_2 \rightarrow 2H_2O}$ represents hydrogen and oxygen forming water. The arrow means “yields” or “forms.” A plus sign separates substances that react together or are produced together.

Letters in parentheses indicate physical states: $\mathrm{(s)}$ means solid, $\mathrm{(l)}$ means liquid, $\mathrm{(g)}$ means gas, and $\mathrm{(aq)}$ means dissolved in water. For example, $\mathrm{AgCl(s)}$ indicates solid silver chloride.

Conservation of matter and balancing

The law of conservation of matter states that the total amount of matter remains constant during a chemical reaction. Because atoms are rearranged rather than lost or gained, a balanced equation has the same number of atoms of each element on both sides. In a closed system, the mass of the reactants therefore equals the mass of the products.

To balance an equation, adjust the coefficients—the numbers placed before chemical formulas. Each coefficient multiplies every atom in the formula that follows it. Do not change a formula’s subscripts: subscripts define the substance, so changing one changes the reaction’s ingredients.

A reliable balancing method

1. Write the correct formulas for the reactants and products.

2. Count the atoms of each element on both sides.

3. Add coefficients until the counts match. Balance elements that occur in fewer formulas first; hydrogen and oxygen are often convenient to balance later.

4. Recount every element and reduce the coefficients to the smallest whole-number ratio.

For example, balance $\mathrm{Al + O_2 \rightarrow Al_2O_3}$. The product contains $2$ aluminum atoms and $3$ oxygen atoms. Use $\mathrm{2Al_2O_3}$ to make the oxygen count $6$, then place $\mathrm{3}$ before $\mathrm{O_2}$ to give $6$ oxygen atoms on the reactant side. The products now contain $4$ aluminum atoms, so place $\mathrm{4}$ before $\mathrm{Al}$.

The balanced equation is:

Each side has $4$ aluminum atoms and $6$ oxygen atoms. The coefficients show relative numbers of particles—or, at the laboratory scale, relative amounts in moles—not equal masses of each substance.

Common reaction patterns

Reaction patterns help organize and predict chemical changes. They are useful guides, but a proposed product must also be chemically possible under the reaction conditions. Recognizing a pattern does not replace balancing: every equation still needs an atom-count check.

• Synthesis: Simpler substances combine, following $\mathrm{A + B \rightarrow AB}$. For example, $\mathrm{2Mg + O_2 \rightarrow 2MgO}$.

• Decomposition: One substance breaks apart, following $\mathrm{AB \rightarrow A + B}$. For example, $\mathrm{2H_2O_2 \rightarrow 2H_2O + O_2}$.

• Single replacement: One element replaces another in a compound, following $\mathrm{A + BC \rightarrow AC + B}$. For example, $\mathrm{Zn + 2HCl \rightarrow ZnCl_2 + H_2}$. Here, zinc replaces hydrogen in hydrochloric acid, producing hydrogen gas.

• Double replacement: Ions in two compounds exchange partners, following $\mathrm{AB + CD \rightarrow AD + CB}$. For example, $\mathrm{AgNO_3 + NaCl \rightarrow AgCl(s) + NaNO_3}$. The $\mathrm{AgCl(s)}$ forms as an insoluble solid called a precipitate.

• Combustion: A substance reacts with oxygen. Complete combustion of a hydrocarbon produces carbon dioxide and water, following $\mathrm{hydrocarbon + O_2 \rightarrow CO_2 + H_2O}$. For example, $\mathrm{CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O}$.

Practice balancing an equation

Balance $\mathrm{N_2 + H_2 \rightarrow NH_3}$ by adjusting coefficients:

Both sides contain $2$ nitrogen atoms and $6$ hydrogen atoms.