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A study on 5 key experiments on chemical reactions

Background Info:

Chemical reactions are an everyday occurrence, but not everyone may realize that. There are six chemical reactions happening all around us, but this experiment will only focus on five. The different types of reactions are: Single displacement/replacement, double displacement/replacement, synthesis, combustion, and decomposition. Some reactions happen to form new compounds, while others may breakdown a previous compound.

In order to find out what elements will displace other elements, it is crucial to look at the activity series of metals. This chart will help you to determine which metals and halogens are the most or least reactive. The most active elements will replace the least active elements in both single and double displacement. However, some metals are not reactive. In this instance, no chemical reaction will occur and that is often indicated by writing “NR” (No Reaction).

Single and double displacement are reactions that form new compounds. Single replacement happens when one element replaces the atoms of another element in a compound. An example of this is:

Cu(s) + 2AgNO3(aq) ? 2Ag(s) + Cu(NO3)2 (aq)

In this equation, copper in its solid form and silver nitrate in a solution produce a silver precipitate and copper(II) nitrate in a solution.

Double displacement is very similar to single displacement in that there is an exchange of ions between two compounds. An example of this is:

Ca(OH)2(aq) + 2HCl(aq) ? CaCl2(aq) + 2H2O (l)

In this equation, a solution of calcium hydroxide plus hydrogen chloride produces calcium chloride in a solution and water.

A huge characteristic of double displacement is the product that is formed from the chemical equation. The products will produce either a precipitate, gas, or water.

Synthesis is a complete different type of reaction where two or more substances react and form a single product. When there are only two elements, the reaction will always be from synthesis. An example of this is:

2Na(s) + Cl2(g) ? 2NaCl(s)

In this equation, sodium in its solid form plus chlorine gas react and produce the compound sodium chloride as a solid.

Decomposition reactions are opposite to synthesis reactions in that decomposition has one compound that breaks down into two or more elements. decompositions normally require an energy source like heat or electricity. An example of this is:

NH4NO3(s) ? N2O(g) + 2H2O(g)

In this reaction, the solid ammonium nitrate breaks down to produce water and nitrous oxide in their gas forms.

Combustion is the last reaction in this experiment and is one of the easiest to classify. This is because the products of combustion are always water and carbon dioxide. In combustion reactions, oxygen combines with another substance, usually hydrocarbons, to release energy in the form of heat. An example of this is:

CH4(g) + 2O2(g) ? CO2(g) + 2H2O(g)

In this equation, the gases methane and oxygen are added together to produce carbon dioxide and water.


To demonstrate and identify different types of reactions.


  1. If the products of a reaction are CO2 and H2O, then combustion has occurred.
  2. If two or more substances combine into one compound, then synthesis has occurred.
  3. If one compound separates into two or more elements and/or compounds, then decomposition has occurred.
  4. If the anions or cations in one compound switch places, then single displacement has occurred.
  5. If the anions or cations in two compounds switch places, then double displacement has occurred.


Experiment #1

The starter produced very small sparks which lit the CH4(g). The methane coming from the gas line was flammable and showed when it produced a flame. The flame was an indication that a combustion occurred and produced H2O and CO2. This experiment was limited by the amount of CH4 present.

Experiment #2

After initial observations, the solution was a clear liquid. However, after heating the solution, there was fizzing and bubbling. This fizzing was probably caused by CO2 being released. After waiting for the solution to evaporate, a new substance was left. This substance was milky white and was slightly powdery. The powder produced a small vinegar smell. this experiment was classified as a Decomposition because the original solution broke down into separate elements. The element that was left appeared to be CaO, or Lime.

Experiment #3

The Fe(s) did not react to the fire. This means that a combustion did not occur. Even after burning the whole skewer, the Fe(s) still did not cause a reaction. Fe is reactive and will cause a synthesis eventually when rust is formed. This experiment was not enough to cause this reaction to take place because enough time didn’t pass.

Experiment #4

The first solution of Mg was a clear, odorless liquid. Once the Zn strip was placed in the solution of Mg, bubbles began forming around the strip. The Mg slowly began depositing on the Zn strip and the original silver color began to turn darker. The Zn strip also had a solid form along the outside. This experiment would be classified as a single replacement

Experiment #5

The NaOH was a clear liquid, while the CuCl2 had a blue tint to it. Once the CuCl2 was added to the solution of NaOH, the liquids mixed and formed a precipitate. This solution instantly formed into a blue-green precipitate with no smell. The liquids turning into a precipitate is a tell-tale sign of the experiment being a double replacement.

Final Thoughts:

If these experiments were to be conducted again, some changes could be implemented to make sure the optimal results happen. For example, in experiment #3, there was no reaction. Adding a flame to Fe(s) should have produced at least a color change. The experiment conducted had absolutely no reaction. Fe(s) should have formed rust through synthesis after enough time had passed. This means that if the experiment were to be conducted again, the flame should be held under Fe(s) for a much longer period of time so that a visible reaction would occur.

Experiment #1 had a major limitation. This experiment was limited by the amount of CH4 in total and the amount of CH4 being used at one time. If the gas knob was turned on full to allow the maximum amount of CH4, the flame was very high. However, if the gas knob was barely turned on and restricted the amount of CH4, the flame was very low.

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