STOY-KEE-OHM-EH-TREE. Big word. Simple idea.
# pieces before the reaction = # pieces after the reaction
We counts pieces whether as (s) solids, (l) liquids or (g) gases.
Here are 4 different, unrelated, chemical reactions.
Left side: reactants Right side: products
Methane is used as a fuel for ovens, homes, water heaters. It combines with oxygen in a combustion reaction to form release heat (fire!), producing water and carbon dioxide.
In terms of what molecules are involved, the diagram below is accurate.
But look closely, and count: What is inaccurate about this image?
Keeping track of the number of atoms is important.
Let’s burn methane gas!
So far so good, in terms of the reactants (ingredients) and the products( stuff that we made)
But is this scale balanced? If so, why? If not, why not?
Look closely. What did we need to do to balance the scale?
“Modern Chemistry” Houghton-Mifflin, Chapter 9
Modern Chemistry textbook online
Much of our knowledge of chemistry is based on the careful quantitative analysis of substances involved in chemical reactions. Composition stoichiometry deals with the mass relationships of elements in compounds. Reaction stoichiometry involves the mass relationships between reactants and products in a chemical reaction.
Reaction stoichiometry, the subject of this chapter, is based on chemical equations and the law of conservation of mass. All reaction stoichiometry calculations start with a balanced chemical equation. This equation gives the relative numbers of moles of reactants and products.
People throughout history have transformed substances by burning them in air. Yet at the dawn of the scientific revolution, very little was known about the process of combustion. In attempting to explain this common phenomenon, chemists of the 18th century developed one of the first universally accepted theories in their field. But as one man would show, scientific theories do not always stand the test of time
Shunning the ancient Greek approach of logical argument based on untested premises, investigators of the 17th century began to understand the laws of nature by observing, measuring, and performing experiments on the world around them. However, this scientific method was incorporated into chemistry slowly. Although early chemists experimented extensively, most considered measurement to be unimportant. This viewpoint hindered the progress of chemistry for nearly a century.
By 1700, combustion was assumed to be the decomposition of a material into simpler substances. People saw burning substances emitting smoke and energy as heat and light. To account for this emission, scientists proposed a theory that combustion depended on the emission of a substance called phlogiston (Greek, meaning “burned”), which appeared
as a combination of energy as heat and light while the material was burning but which could not be detected beforehand. The phlogiston theory was used to explain many chemical observations of the day. For example, a lit candle under a glass jar burned until the surrounding air became saturated with phlogiston, at which time the flame died because the air inside could not absorb more phlogiston.
By the 1770s, the phlogiston theory had gained universal acceptance. At that time, chemists also began to experiment with air, which was generally believed to be an element. In 1772, when Daniel Rutherford found that a mouse kept in a closed container soon died, he explained the results based on the phlogiston theory. Like a burning candle, the mouse emitted phlogiston; when the air could hold no more phlogiston, the mouse died. Thus, Rutherford figured that the air in the container had become “phlogisticated air.”
A couple of years later, Joseph Priestley obtained a reddish powder when he heated mercury in the air. He assumed that the powder was mercury devoid of phlogiston. But when he heated the powder, an unexpected result occurred: Metallic mercury, along with a gas that allowed a candle to burn, formed. Following the phlogiston theory, he believed this gas that supports combustion to be “dephlogisticated air.”
People throughout history have transformed substances by burning them in air. Yet at the dawn of the scientific revolution, very little was known about the process of combustion. In attempting to explain this common phenomenon, chemists of the 18th century developed one of the first universally accepted theories in their field. But as one man would show, scientific theories do not always stand the test of time. Antoine Laurent Lavoisier was a meticulous scientist. He realized that Rutherford and Priestley had carefully observed and described their experiments but had not measured the mass of anything.
Unlike his colleagues, Lavoisier knew the importance of using a balance. He measured the masses of reactants and products and compared them. He observed that the total mass of the reactants equaled the total mass of the products. Based on these observations, which supported what would become known as the law of conservation of mass , Lavoisier endeavored to explain the results of Rutherford and Priestley.
video clip: Conservation of Mass, from Einstein’s Big Idea – Nature is a closed system
Lavoisier Extended video clip – 13 minutes