Law of Conservation of Mass & Definite Composition

 This is part of Year 11 HSC Chemistry course under the topic of Chemical Reactions and Stoichiometry

HSC Chemistry Syllabus

  • Relate stoichiometry to the law of conservation of mass in chemical reactions by investigating:
    • Balancing chemical equations (ACSCH039)
    • Solving problems regarding mass changes in chemical reactions (ACSCH046)

    Law of Conservation of Mass

    The principle of the conservation of mass is rooted in Antoine Lavoisier's 1789 discovery, establishing that mass is neither created nor destroyed in chemical reactions but merely undergoes transformation. This foundational principle is pivotal in quantitative chemistry, particularly in balancing chemical equations. In essence, the total mass remains constant before and after any chemical reaction.


    Consider the combustion of fuel as a practical demonstration of the law of conservation of mass. In our study's fifth stage, we dissect combustion reactions. These reactions encompass molecular oxygen, culminating in the emergence of carbon dioxide, water, and heat. Upon combusting a hydrocarbon fuel like octane (\(C_8H_{18}\)), one might perceive the fuel as vanished. Contrarily, the fuel metamorphoses into water vapor and carbon dioxide, which eludes our naked eye, seemingly erasing the fuel.


    This tenet similarly applies to precipitation reactions, such as when forming an insoluble salt like lead sulfate.


    \[CoSO_4(aq) + Pb(NO_3)_2(aq) \rightarrow PbSO_4(s)\]








    No. molecules on left






    No. molecules on right







    Definite Composition


    Regardless of its source, be it from your salt shaker or the sea, sodium chloride retains its composition. This uniformity is articulated by the law of definite composition: a chemical compound invariably maintains a consistent elemental ratio, irrespective of its source or preparation method.


    The mass fraction of an element in a compound is the fraction that element's mass contributes to the compound. Its calculation involves dividing the element's mass by the compound's overall mass. The mass percent is simply this fraction, expressed as a percentage.


    $$\text{Mass Fraction} = \frac{\text{mass of element X in compound A}}{\text{mass of compound A}}$$


    $$\text{Mass percent} = \text{mass fraction} \times 100 $$


    For an illustration, consider 18g of fruit with 3 apples, each weighing 3g. The mass fraction of apples is:


    $$\text{Mass Fraction} = \frac{3 \, \text{apples} \times 3 \, \text{g/apple}}{18 \, \text{g of fruit}} $$

    $$\text{Mass percent}  = 0.5 \times 100 = 50\% $$


    Therefore, apples constitute `50\%` of the fruit's mass, with a mass fraction of 0.5.


    In a similar manner, each element in a compound has an unchanging mass fraction and mass percent. Consider calcium carbonate, found abundantly in seashells, marble, and coral. Comprising calcium, oxygen, and carbon, a mass analysis of a 20g sample yields:


    Comparing elemental occurrences on both sides:



    Mass fraction

    Mass percent

    8g calcium

    0.4 calcium

    40% calcium

    2.4g carbon

    0.12 carbon

    12% carbon

    9.6g oxygen

    0.48 oxygen

    48% oxygen





    Regardless of the sample size, the element's mass is contingent on the total sample mass due to the constant mass fraction---an embodiment of the law of definite composition.

    Measuring Liquids in Chemical Reactions

    The density of a liquid is commonly expressed in the unit \( \text{g/mL} \). Given both the density and volume of a liquid, its mass can be determined using the relation:

    $$\text{Mass} = \text{volume} \times \text{density}$$

    Example: Calculate the mass of 30 mL of methanol, given its density as \(0.79 \, \text{g/mL}\).

    $$\text{Mass} = \text{volume} \times \text{density} $$
    $$\text{Mass} = 30 \, \text{mL} \times 0.79 \, \text{g/mL} $$
    $$\text{Mass} = 23.7 \, \text{g}$$