Chemical Reactions and the Law of Conservation of Mass: Seventh Grade Science
Chemical reactions rearrange atoms into new molecular configurations without creating or destroying matter, in strict accordance with the Law of Conservation of Mass.
Antoine Lavoisier and the Conservation of Mass
In 1789, French chemist Antoine-Laurent de Lavoisier established modern quantitative chemistry by formulating the Law of Conservation of Mass. Through experiments using sealed glass vessels (closed systems), Lavoisier demonstrated that in any chemical reaction, the total mass of the products must equal the total mass of the reactants. Matter can neither be created nor destroyed by chemical reactions or physical transformations.
+-----------------------------------------------------------------------------+
| LAW OF CONSERVATION OF MASS |
| |
| TOTAL MASS OF REACTANTS === TOTAL MASS OF PRODUCTS |
| |
| Closed Vessel Experiment: |
| Reactants: 100.0 g Vinegar + 10.0 g Baking Soda = 110.0 g Initial Mass |
| Chemical Reaction Occurs: Effervescent bubbling of CO2 gas |
| Products: 105.6 g Sodium Acetate Sol. + 4.4 g CO2 Gas = 110.0 g Mass |
| |
| Result: Mass is conserved perfectly in a sealed container! |
+-----------------------------------------------------------------------------+
The Atomic Basis of Conservation
Lavoisier's macro-scale observations were explained at the atomic scale by John Dalton's atomic theory: chemical reactions are the breaking, rearrangement, and reforming of chemical bonds between atoms. Because atoms are indestructible particles that cannot be created from nothing or destroyed in chemical reactions, every single atom present in the reactant molecules must be accounted for among the product molecules. Atoms are rearranged into new combinations, preserving total mass.
The Anatomy of a Chemical Equation
A chemical equation is a symbolic mathematical model that represents a chemical reaction using chemical formulas, stoichiometric coefficients, and state symbols:
+-----------------------------------------------------------------------------+
| ANATOMY OF A CHEMICAL EQUATION |
| |
| REACTANTS PRODUCTS |
| 2 H2(g) + O2(g) --------> 2 H2O(g) |
| ^ ^ ^ ^ ^ |
| | | | | | |
| Coefficient Subscript Coefficient Subscript |
| (Molecules) (Atoms per molecule) (Molecules) |
| |
| Atom Inventory: |
| - Reactants: (2 * 2) = 4 Hydrogen atoms; (1 * 2) = 2 Oxygen atoms |
| - Products: (2 * 2) = 4 Hydrogen atoms; (2 * 1) = 2 Oxygen atoms |
| Equation is Balanced! |
+-----------------------------------------------------------------------------+
- Reactants: The starting chemical substances undergoing the reaction, written on the left side of the reaction arrow.
- Products: The newly synthesized chemical substances produced by the reaction, written on the right side of the reaction arrow.
- Reaction Arrow (-------->): Represents the chemical transformation, read as "yields" or "produces."
- Subscripts: Small numbers written to the lower right of an elemental symbol that indicate the exact number of atoms of that element bonded within one molecule or formula unit (e.g., in H2O, the subscript 2 indicates two hydrogen atoms bonded to one oxygen atom). Subscripts can never be changed when balancing an equation, as doing so alters the chemical identity of the substance.
- Stoichiometric Coefficients: Full-sized numbers placed in front of chemical formulas that indicate the relative number of molecules or moles of that substance participating in the reaction. Coefficients multiply all atoms within the formula.
Balancing Chemical Equations: Step-by-Step Methodology
An unbalanced chemical equation violates the Law of Conservation of Mass by showing unequal numbers of atoms on opposite sides of the reaction arrow. Balancing an equation requires adjusting coefficients until the number of atoms of each element is equal on both sides:
+-----------------------------------------------------------------------------+
| BALANCING COMBUSTION OF PROPANE |
| |
| Unbalanced: C3H8 + O2 --------> CO2 + H2O |
| |
| Step 1: Balance Carbon (C) atoms: |
| - Reactants have 3 Carbons; add coefficient 3 to CO2: |
| C3H8 + O2 --------> 3 CO2 + H2O |
| |
| Step 2: Balance Hydrogen (H) atoms: |
| - Reactants have 8 Hydrogens; add coefficient 4 to H2O: |
| C3H8 + O2 --------> 3 CO2 + 4 H2O |
| |
| Step 3: Balance Oxygen (O) atoms: |
| - Products have (3 * 2) + (4 * 1) = 10 Oxygen atoms |
| - Reactants have O2; add coefficient 5 to O2: |
| C3H8 + 5 O2 --------> 3 CO2 + 4 H2O |
| |
| Balanced Check: 3 C, 8 H, 10 O on both sides! |
+-----------------------------------------------------------------------------+
Closed Versus Open Systems
Apparent violations of the Law of Conservation of Mass stem from a failure to account for all matter entering or escaping open systems:
- Open System: Matter and thermal energy can be exchanged freely between the system and its surrounding environment. For example, burning a 5.0-kilogram wooden log in an open campfire leaves behind only 0.2 kilograms of solid ash. A novice might conclude that 4.8 kilograms of matter was destroyed. However, the missing mass escaped into the surrounding atmosphere as gaseous carbon dioxide, water vapor, and carbon monoxide.
- Closed System: Energy can enter or leave, but matter is physically prevented from crossing the system boundary. If the identical 5.0-kilogram wooden log were sealed inside a steel chamber with sufficient oxygen and ignited, the total mass of the sealed chamber and its contents would remain unchanged before, during, and after combustion.
Exercises and Solutions
Exercise 1: Balancing Chemical Equations and Verifying Mass Conservation. The chemical process of industrial ammonia synthesis (the Haber-Bosch process) reacts diatomic nitrogen gas (N2) with diatomic hydrogen gas (H2) to produce ammonia gas (NH3). (a) Write the unbalanced chemical equation; (b) Balance the chemical equation by inserting the correct stoichiometric coefficients; (c) If 28.0 grams of nitrogen gas react completely with exactly 6.0 grams of hydrogen gas, calculate the mass of ammonia gas produced, explaining your calculation using the Law of Conservation of Mass.
Solution 1: (a) The unbalanced chemical equation is: N2 + H2 --------> NH3. (b) To balance the equation: The reactant side contains two nitrogen atoms (N2), so place a coefficient of 2 in front of NH3 on the product side: N2 + H2 --------> 2 NH3. This gives 2 * 3 = 6 hydrogen atoms on the product side. To match this, place a coefficient of 3 in front of H2 on the reactant side: N2 + 3 H2 --------> 2 NH3. Checking the inventory confirms: 2 Nitrogen atoms and 6 Hydrogen atoms on both sides of the equation. (c) According to the Law of Conservation of Mass, the total mass of the products must equal the total mass of the consumed reactants. Because 28.0 grams of nitrogen react completely with 6.0 grams of hydrogen: Total Mass of Products = Mass of Nitrogen + Mass of Hydrogen = 28.0 g + 6.0 g = 34.0 grams of ammonia gas (NH3) produced.
Exercise 2: Diagnosing Apparent Mass Violations in Open Versus Closed Reactions. A student conducts two separate chemical demonstrations on an open electronic balance. In Demonstration 1, the student places an open flask containing 50.0 grams of hydrochloric acid (HCl) on the balance, tares it, drops in a 2.0-gram strip of solid magnesium ribbon, and watches effervescent bubbling. When bubbling stops, the balance reads 51.8 grams instead of 52.0 grams. In Demonstration 2, the student places 10.0 grams of steel wool (pure iron) on the balance and ignites it with a battery. As the steel wool glows and oxidizes into iron oxide, the balance reading increases from 10.0 grams to 13.8 grams. Explain why Demonstration 1 appeared to lose mass while Demonstration 2 appeared to gain mass, proving that neither reaction violated the Law of Conservation of Mass.
Solution 2: Neither demonstration violates the Law of Conservation of Mass; both were conducted in open systems where gaseous matter crossed the boundary undetected by the balance. In Demonstration 1, magnesium reacts with hydrochloric acid to produce aqueous magnesium chloride and hydrogen gas: Mg + 2 HCl --------> MgCl2 + H2. Because the reaction occurred in an open flask, the low-density hydrogen gas (H2) escaped into the atmosphere. The 0.2 grams of missing mass represents the exact mass of hydrogen gas that escaped. If the flask had been sealed with an airtight stopper, the mass would have remained 52.0 grams. In Demonstration 2, iron metal reacts with atmospheric oxygen gas at high temperature to form solid iron(III) oxide: 4 Fe + 3 O2 --------> 2 Fe2O3. The balance initially weighed only the solid iron steel wool (10.0 grams). During combustion, iron atoms bonded chemically with oxygen atoms drawn directly from the surrounding air. The 3.8-gram mass increase represents the mass of oxygen gas pulled from the atmosphere into the solid chemical structure. In a closed container, the initial mass of the iron plus the enclosed oxygen gas would equal the final mass of the iron oxide, confirming conservation of mass.