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Conservation of matter

Module by: Free High School Science Texts Project. E-mail the author

Energy changes in chemical reactions

All reactions involve some change in energy. During a physical change in matter, such as the evaporation of liquid water to water vapour, the energy of the water molecules increases. However, the change in energy is much smaller than in chemical reactions.

When a chemical reaction occurs, some bonds will break, while new bonds may form. Energy changes in chemical reactions result from the breaking and forming of bonds. For bonds to break, energy must be absorbed. When new bonds form, energy will be released because the new product has a lower energy than the `in between' stage of the reaction when the bonds in the reactants have just been broken.

In some reactions, the energy that must be absorbed to break the bonds in the reactants is less than the total energy that is released when new bonds are formed. This means that in the overall reaction, energy is released. This type of reaction is known as an exothermic reaction. In other reactions, the energy that must be absorbed to break the bonds in the reactants is more than the total energy that is released when new bonds are formed. This means that in the overall reaction, energy must be absorbed from the surroundings. This type of reaction is known as an endothermic reaction. Most decomposition reactions are endothermic and heating is needed for the reaction to occur. Most synthesis reactions are exothermic, meaning that energy is given off in the form of heat or light.

More simply, we can describe the energy changes that take place during a chemical reaction as:

Total energy absorbed to break bonds - Total energy released when new bonds form

So, for example, in the reaction...

2Mg+O2 → 2MgO2Mg+O2→ 2MgO

Energy is needed to break the O-OO-O bonds in the oxygen molecule so that new Mg-OMg-O bonds can be formed, and energy is released when the product (MgOMgO) forms.

Despite all the energy changes that seem to take place during reactions, it is important to remember that energy cannot be created or destroyed. Energy that enters a system will have come from the surrounding environment and energy that leaves a system will again become part of that environment. This is known as the conservation of energy principle.

Definition 1: Conservation of energy principle
Energy cannot be created or destroyed. It can only be changed from one form to another.

Chemical reactions may produce some very visible and often violent changes. An explosion, for example, is a sudden increase in volume and release of energy when high temperatures are generated and gases are released. For example, NH4NO3NH4NO3 can be heated to generate nitrous oxide. Under these conditions, it is highly sensitive and can detonate easily in an explosive exothermic reaction.

Conservation of atoms and mass in reactions

The total mass of all the substances taking part in a chemical reaction is conserved during a chemical reaction. This is known as the law of conservation of mass. The total number of atoms of each element also remains the same during a reaction, although these may be arranged differently in the products.

We will use two of our earlier examples of chemical reactions to demonstrate this:

1. The decomposition of hydrogen peroxide into water and oxygen

2 H 2 O 2 → 2 H2O + O 2 2 H 2 O 2 → 2 H2O + O 2

Figure 1
Figure 1 (CG10C4_007.png)

Left hand side of the equation

Total atomic mass=(4×1)+(4×16)=68uTotal atomic mass=(4×1)+(4×16)=68u

Number of atoms of each element=(4×H)+(4×O)Number of atoms of each element=(4×H)+(4×O)

Right hand side of the equation

Total atomic mass=(4×1)+(4×16)=68uTotal atomic mass=(4×1)+(4×16)=68u

Number of atoms of each element=(4×H)+(4×O)Number of atoms of each element=(4×H)+(4×O)

Both the atomic mass and the number of atoms of each element are conserved in the reaction.

2. The synthesis of magnesium and oxygen to form magnesium oxide

2Mg+O2 → 2MgO2Mg+O2→ 2MgO
(1)

Figure 2
Figure 2 (CG10C4_008.png)

Left hand side of the equation

Total atomic mass=(2×24,3)+(2×16)=80,6uTotal atomic mass=(2×24,3)+(2×16)=80,6u

Number of atoms of each element=(2×Mg)+(2×O)Number of atoms of each element=(2×Mg)+(2×O)

Right hand side of the equation

Total atomic mass=(2×24,3)+(2×16)=80,6uTotal atomic mass=(2×24,3)+(2×16)=80,6u

Number of atoms of each element=(2×Mg)+(2×O)Number of atoms of each element=(2×Mg)+(2×O)

Both the atomic mass and the number of atoms of each element are conserved in the reaction.

Activity : The conservation of atoms in chemical reactions

Materials:

  1. Coloured marbles or small balls to represent atoms. Each colour will represent a different element.
  2. Prestik

Method:

  1. Choose a reaction from any that have been used in this chapter or any other balanced chemical reaction that you can think of. To help to explain this activity, we will use the decomposition reaction of calcium carbonate to produce carbon dioxide and calcium oxide. CaCO3→CO2+CaOCaCO3→CO2+CaO
  2. Stick marbles together to represent the reactants and put these on one side of your table. In this example you may for example join one red marble (calcium), one green marble (carbon) and three yellow marbles (oxygen) together to form the molecule calcium carbonate (CaCO3CaCO3).
  3. Leaving your reactants on the table, use marbles to make the product molecules and place these on the other side of the table.
  4. Now count the number of atoms on each side of the table. What do you notice?
  5. Observe whether there is any difference between the molecules in the reactants and the molecules in the products.

Discussion

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 molecules in the reactants and products is not the same. The arrangement of atoms is not conserved during the reaction.

Experiment: Conservation of matter

Aim:

To prove the law of conservation of matter experimentally.

Materials:

Test tubes; glass beaker; lead (II) nitrate; sodium iodide; hydrochloric acid; bromothymol blue; Cal-C-Vita tablet, plastic bag; rubber band; mass meter

Warning:

Always be careful when handling chemicals (particularly strong acids like hydrochloric acid) as you can burn yourself badly.

Method:

Reaction 1

  1. Carefully weigh out 5 g of lead (II) nitrate.
  2. Dissolve the lead nitrate in 100 ml of water.
  3. Weigh the lead nitrate solution.
  4. Weigh out 4,5 g of sodium iodide and dissolve this in the lead (II) nitrate solution.
  5. Weigh the beaker containing the lead nitrate and sodium iodide mixture.
Reaction 2
  1. Measure out 20 ml of sodium hydroxide.
  2. Add a few drops of bromothymol blue to the sodium hydroxide.
  3. Weigh the sodium hydroxide.
  4. Weigh 5 ml of hydrochloric acid.
  5. Add 5 ml of hydrochloric acid to the sodium hydroxide. Repeat this step until you observe a colour change (this should occur around 20 ml).
  6. Weigh the final solution.
Reaction 3
  1. Measure out 100 ml of water into a beaker.
  2. Weigh the beaker with water in it.
  3. Place the Cal-C-Vita tablet into the plastic bag.
  4. Weigh the Cal-C-Vita tablet and the plastic bag.
  5. Place the plastic bag over the beaker, being careful to not let the tablet fall into the water
  6. Seal the bag around the beaker using the rubber band. Drop the tablet into the water.
  7. Observe what happens.
  8. Weigh the bag and beaker containing the solution.

Results:

Fill in the following table for reactants (starting materials) and products (ending materials) masses. For the second reaction, you will simply take the mass of 5 ml of hydrochloric acid and multiply it by how many amounts you put in, for example, if you put 4 amounts in, then you would have 20 ml and 4 times the mass of 5 ml.

Table 1
  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 mass at the start of the reaction is the same as the mass at the end of the reaction. You may have found that these masses differed slightly, but this is due to errors in measurements and in performing experiments (all scientists make some errors in performing experiments).

Law of constant composition

In any given chemical compound, the elements always combine in the same proportion with each other. This is the law of constant proportion.

The law of constant composition says that, in any particular chemical compound, all samples of that compound will be made up of the same elements in the same proportion or ratio. For example, any water molecule is always made up of two hydrogen atoms and one oxygen atom in a 2:1 ratio. If we look at the relative masses of oxygen and hydrogen in a water molecule, we see that 94% of the mass of a water molecule is accounted for by oxygen and the remaining 6% is the mass of hydrogen. This mass proportion will be the same for any water molecule.

This does not mean that hydrogen and oxygen always combine in a 2:1 ratio to form H2OH2O. Multiple proportions are possible. For example, hydrogen and oxygen may combine in different proportions to form H2O2H2O2 rather than H2OH2O. In H2O2H2O2, the H:O ratio is 1:1 and the mass ratio of hydrogen to oxygen is 1:16. This will be the same for any molecule of hydrogen peroxide.

Volume relationships in gases

In a chemical reaction between gases, the relative volumes of the gases in the reaction are present in a ratio of small whole numbers if all the gases are at the same temperature and pressure. This relationship is also known as Gay-Lussac's Law.

For example, in the reaction between hydrogen and oxygen to produce water, two volumes of H2H2 react with 1 volume of O2O2 to produce 2 volumes of H2OH2O.

2H2+O2→2H2O2H2+O2→2H2O

In the reaction to produce ammonia, one volume of nitrogen gas reacts with three volumes of hydrogen gas to produce two volumes of ammonia gas.

N2+3H2→2NH3N2+3H2→2NH3

This relationship will also be true for all other chemical reactions.

Summary

The following video provides a summary of the concepts covered in this chapter.

Figure 3
Physical and chemical change

  1. Matter does not stay the same. It may undergo physical or chemical changes.
  2. A physical change means that the form of matter may change, but not its identity. For example, when water evaporates, the energy and the arrangement of water molecules will change, but not the structure of the water molecules themselves.
  3. During a physical change, the arrangement of particles may change but the mass, number of atoms and number of molecules will stay the same.
  4. Physical changes involve small changes in energy and are easily reversible.
  5. A chemical change occurs when one or more substances change into other materials. A chemical reaction involves the formation of new substances with different properties. For example, magnesium and oxygen react to form magnesium oxide (MgOMgO)
  6. A chemical change may involve a decomposition or synthesis reaction. During chemical change, the mass and number of atoms is conserved, but the number of molecules is not always the same.
  7. Chemical reactions involve larger changes in energy. During a reaction, energy is needed to break bonds in the reactants and energy is released when new products form. If the energy released is greater than the energy absorbed, then the reaction is exothermic. If the energy released is less than the energy absorbed, then the reaction is endothermic. Chemical reactions are not easily reversible.
  8. Decomposition reactions are usually endothermic and synthesis reactions are usually exothermic.
  9. The law of conservation of mass states that the total mass of all the substances taking part in a chemical reaction is conserved and the number of atoms of each element in the reaction does not change when a new product is formed.
  10. The conservation of energy principle states that energy cannot be created or destroyed, it can only change from one form to another.
  11. The law of constant composition states that in any particular compound, all samples of that compound will be made up of the same elements in the same proportion or ratio.
  12. Gay-Lussac's Law states that in a chemical reaction between gases, the relative volumes of the gases in the reaction are present in a ratio of small whole numbers if all the gases are at the same temperature and pressure.

End of chapter exercises

  1. For each of the following definitions give one word or term:
    1. A change that can be seen or felt, where the particles involved are not broken up in any way
    2. The formation of new substances in a chemical reaction
    3. A reaction where a new product is formed from elements or smaller compounds
    Click here for the solution
  2. State the conservation of energy principle.
    Click here for the solution
  3. Explain how a chemical change differs from a physical change.
    Click here for the solution
  4. Complete the following table by saying whether each of the descriptions is an example of a physical or chemical change:
    Table 2
    DescriptionPhysical or chemical
    hot and cold water mix together 
    milk turns sour 
    a car starts to rust 
    food digests in the stomach 
    alcohol disappears when it is placed on your skin 
    warming food in a microwave 
    separating sand and gravel 
    fireworks exploding 
    Click here for the solution
  5. For each of the following reactions, say whether it is an example of a synthesis or decomposition reaction:
    1. ( NH4 )2CO3 → NH 3+CO 2+H2O ( NH4 )2CO3 → NH 3+CO 2+H2O
    2. N2(g)+3H2(g)→2NH3N2(g)+3H2(g)→2NH3
    3. CaCO3(s)→CaO+CO2CaCO3(s)→CaO+CO2
    Click here for the solution
  6. For the following equation: CaCO3(s)→CaO+CO2CaCO3(s)→CaO+CO2 show that the 'law of conservation of mass' applies.
     Click here for the solution

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