Although life on Earth may seem rather stable and unchanging—the tide goes in and out, the Sun rises and sets, and the months bleed on the same as always—in the grand scheme of things, our universe is actually a rather dynamic place.
Everyday, millions of stars are born and die, and in the end, the same thing will happen to our own Sun.
A few billion years from now, as our star begins to transition into a red giant, temperatures on our world will increase, and life will be extinguished. Just a few billion years after that, once the Sun exhausts its supply of material to sustain nuclear fusion, it will begin its death throes. When this happens, it will cast off its outer layers and, eventually, fade into darkness.
Nothing in our universe is eternal...or is it?
Atoms are the building blocks of matter. They, quite literally, make our universe what it is. When we die, our bodies do not turn into nothing; rather, they are broken down into their constituent parts and recycled into the ecosystem. In short, our atoms go on long after we are gone.
But just how long can atoms last? Will they eventually be broken down into...nothing?
To answer this question, you need to understand a little bit about how atoms work. As you may know, atoms contain protons and neutrons, and they are surrounded by a "shell" of electrons. The number of electrons that are found in the cloud are equal to the number of protons. This helps create stability.
Ultimately, the number of protons is what determines the atomic number. For example, helium has two protons, so its atomic number is two (and it appears second on the periodic table of elements). The number of neutrons that are found in an atom are generally consistent, but not always. And if an atom doesn't have the "correct" number of neutrons, sometimes, the atom may lose a neutron (kind of like how you lose a sock in the wash). When this happen, the atom becomes unstable and, in an attempt to become a stable atom, it shoots off subatomic particles. Most often, the atom will shoot off an electron.
This is how atoms breakdown.
Anytime that you have a heavy atom, there is some risk that it will spontaneously start to break down into smaller particles. This is known as "radioactive decay." This is just a very basic breakdown. Please see the link for more on radioactive decay.
To return to the point at hand, unfortunately, this is a stochastic process (which means that it has "a random probability distribution, or pattern that may be analyzed statistically, but may not be predicted precisely"). In other words, we can't pinpoint exactly when a breakdown will occur - when a subatomic particle will be shot off; however, since we can analyze the pattern, we can determine how many atoms will decay over an average time, which is called the"half-life," and it is a very reliable estimate.
Since an atom has a finite number of protons and neutrons, it will generally emit particles until it gets to a point where its half-life is so long, it is effectively stable. For example, Bismuth-209 is believed to have the longest decay rate. It undergoes something known as "alpha decay," and it's half-life is over a billion times longer than the current estimated age of the universe.
So for all intents and purposes, Bismuth-209 is basically eternal.
That said, true eternal life depends on whether or not protons can decay. Some scientists have put forth hypotheses related to this, and it is referred to as "proton decay" (a hypothetical form of radioactive decay). According to one idea, the Georgi–Glashow model, protons transition into a positron and a neutral pion, which then decays into 2 gamma ray photons. Estimates put the half-life for protons at 1.29×1034 years.
That, if you don't know, is a super long time; however, there is no experimental evidence to confirm proton decay. But research that is being conducted at some of the world's mega laboratories may, with a bit of luck and hard science, reveal something in the future.
In a chemical reaction the total mass of all the substances taking part in the reaction remains the same. Also, the number of atoms in a reaction remains the same. Mass cannot be created or destroyed in a chemical reaction.
Law of conservation of massThe law of conservation of mass states that the total mass of substances taking part in a chemical reaction is conserved during the reaction.
Table 13.1 illustrates this law for the decomposition of hydrogen peroxide.
We will use the reaction of hydrogen and oxygen to form water in this activity.
Coloured modelling clay rolled into balls or marbles and prestik to represent atoms. Each colour will represent a different element.
-
Build your reactants. Use marbles and prestik or modelling clay to represent the reactants and put these on one side of your table. Make at least ten (\(\text{H}_{2}\)) units and at least five (\(\text{O}_{2}\)) units.
-
Place the \(\text{H}_{2}\) and \(\text{O}_{2}\) units on a table. The table represents the “test tube” where the reaction is going to take place.
-
Now count the number of atoms (\(\text{H}\) and \(\text{O}\)) you have in your “test tube”. Fill in the reactants column in the table below. Refer to Table 13.1 to help you fill in the mass row.
-
Let the reaction take place. Each person can now take the \(\text{H}\) and \(\text{O}\) unit and use them to make water units. Break the \(\text{H}\) and \(\text{O}\) units apart and build \(\text{H}_{2}\text{O}\) units with the parts. These are the products. Place the products on the table.
-
When the “reaction” has finished (i.e. when all the \(\text{H}\) and \(\text{O}\) units have been used) count the number of atoms (\(\text{H}\) and \(\text{O}\)) and complete the table.
-
What do you notice about the number of atoms for the reactants, compared to the products?
-
Write a balanced equation for this reaction and use your models to build this equation.
| Reactants | Products | |
| Number of molecules | ||
| Mass | ||
| Number of atoms |
You should have noticed that the number of atoms in the reactants is the same as the number of atoms in the product. The number of atoms is conserved during the reaction. However, you will also see that the number of molecules in the reactants and products are not the same. The number of molecules is not conserved during the reaction.
To prove the law of conservation of matter experimentally.
Reaction 1:
3 beakers; silver nitrate; sodium iodide; mass meter
Reaction 2:
hydrochloric acid; bromothymol blue; sodium hydroxide solution; mass meter
Reaction 3:
any effervescent tablet (e.g. Cal-C-Vita tablet), balloon; rubber band; mass meter; test tube; beaker
Always be careful when handling chemicals (particularly strong acids like hydrochloric acid) as you can burn yourself badly.
Reaction 1
-
Solution 1: In one of the beakers dissolve \(\text{5}\) \(\text{g}\) of silver nitrate in \(\text{100}\) \(\text{mL}\) of water.
-
Solution 2: In a second beaker, dissolve \(\text{4,5}\) \(\text{g}\) of sodium iodide in \(\text{100}\) \(\text{mL}\) of water.
-
Determine the mass of each of the reactants.
-
Add solution 1 to solution 2. What do you observe? Has a chemical reaction taken place?
-
Determine the mass of the products.
-
What do you notice about the masses?
-
Write a balanced equation for this reaction.
Reaction 2:
-
Solution 1: Dissolve \(\text{0,4}\) \(\text{g}\) of sodium hydroxide in \(\text{100}\) \(\text{mL}\) of water. Add a few drops of bromothymol blue indicator to the solution.
-
Solution 2: Measure \(\text{100}\) \(\text{mL}\) of \(\text{0,1}\) \(\text{mol·dm$^{-3}$}\) hydrochloric acid solution into a second beaker.
-
Determine the mass of the reactants.
-
Add small quantities of solution 2 to solution 1 (you can use a plastic pipette for this) until a colour change has taken place. Has a chemical reaction taken place?
-
Determine the mass of hydrochloric acid added. (You do this by weighing the remaining solution and subtracting this from the starting mass)
-
Compare the mass before the reaction to the total mass after the reaction. What do you notice?
-
Write a balanced equation for this reaction.
Reaction 3
-
Half fill a large test tube with water.
-
Determine the mass of the test tube and water.
-
Break an effervescent tablet in two or three pieces and place them in a balloon.
-
Determine the mass of the balloon and tablet.
-
Fit the balloon tightly to the test tube, being careful to not drop the contents into the water. You can stand the test tube in a beaker to help you do this.
-
Determine the total mass of the test tube and balloon.
-
Lift the balloon so that the tablet goes into the water. What do you observe? Has a chemical reaction taken place?
-
Determine the mass of the test tube balloon combination.
-
What do you observe about the masses before and after the reaction?
Fill in the following table for the total mass of reactants (starting materials) and products (ending materials).
| Reaction 1 | Reaction 2 | Reaction 3 | |
| Reactants | |||
| Products |
Add the masses for the reactants for each reaction. Do the same for the products. For each reaction compare the mass of the reactants to the mass of the products. What do you notice? Is the mass conserved?
In the experiment above you should have found that the total mass at the start of the reaction is the same as the mass at the end of the reaction. Mass does not appear or disappear in chemical reactions. Mass is conserved, in other words, the total mass you start with is the total mass you will end with.
Textbook Exercise 13.2
Complete the following chemical reactions to show that atoms and mass are conserved. For each reaction give the total molecular mass of the reactants and the products.
Hydrogen gas combines with nitrogen gas to form ammonia.
Solution not yet available
Hydrogen peroxide decomposes (breaks down) to form hydrogen and oxygen.
Solution not yet available
Calcium and oxygen gas react to form calcium oxide.
Solution not yet available