Magnesium is a highly reactive, silvery-white metal that readily participates in combustion. When magnesium is ignited, it combines rapidly with oxygen from the surrounding air, producing a brilliant burst of light and significant heat. This energetic transformation is a fundamental example of an oxidation-reduction reaction, where the metal loses electrons and the oxygen gains them. The process of combustion requires a representation that accurately reflects the chemical change, leading to the necessity of a balanced chemical equation.
Reactants and Products
The components that begin this chemical transformation are called the reactants. In this specific combustion, the reactants are solid magnesium (Mg) and oxygen gas (\(\text{O}_2\)). Oxygen exists naturally as a diatomic molecule, meaning two oxygen atoms are chemically bonded together. The sole substance created by this reaction is the product, which is magnesium oxide (MgO). This ionic compound forms when one magnesium atom pairs perfectly with one oxygen atom, resulting in a neutral, fine, white, powdery solid.
Setting Up the Unbalanced Equation
To represent this transformation chemically, the formulas for the reactants and products are arranged into an equation format. The reactants are placed on the left side, separated by a plus sign, and an arrow points toward the product on the right side, indicating the direction of the chemical change. The initial, unbalanced expression is written as \(\text{Mg} + \text{O}_2 \rightarrow \text{MgO}\). This preliminary equation is incomplete because it does not yet account for the number of individual atoms present on each side.
Applying the Law of Conservation of Mass
A chemical equation must adhere to the Law of Conservation of Mass, meaning the total number of atoms for each element must be identical on both the reactant and product sides. The unbalanced equation, \(\text{Mg} + \text{O}_2 \rightarrow \text{MgO}\), violates this law because there are two oxygen atoms on the left but only one on the right. To correct this discrepancy, coefficients (whole numbers placed in front of the formulas) must be introduced. Placing a coefficient of 2 in front of the product, \(\text{MgO}\), successfully balances the oxygen atoms by creating two units of magnesium oxide. The final step requires placing a coefficient of 2 in front of the reactant \(\text{Mg}\) to match the two magnesium atoms now present in the product, satisfying the conservation of mass principle.
The Final Balanced Equation and Visible Results
The final, correctly balanced chemical equation for the combustion of magnesium is \(\text{2Mg} + \text{O}_2 \rightarrow \text{2MgO}\). This equation specifies that two atoms of solid magnesium react with one molecule of gaseous oxygen to yield two formula units of solid magnesium oxide. The coefficients of 2:1:2 represent the stoichiometric ratio, detailing the exact proportions in which the substances react and are produced. The combustion is highly exothermic, releasing a significant amount of energy as intense heat and light. This light has historically been utilized in early photography as flash powder and is still used today in flares and fireworks, producing a plume of white smoke that settles as the fine, powdery residue of magnesium oxide.