Lab: Classification of Chemical Reactions

Introduction

Chemistry is about transformation. Reactants are consumed and products are formed. We take many daily reactions for granted. For example, your body is transforming food by combining it with oxygen to make carbon dioxide and water vapor as a way to generate the energy you need. When we study chemistry we distill matters to simpler, though less familiar, reactions. There are many, many types of chemical reactions. But there are just a few that form the starting point for a student’s understanding of them. Here are some brief descriptions of these introductory reaction types.

Synthesis Reactions

In a synthesis or combination reaction, reactants combine to form a single product. One iconic example is a metal combining with a non-metal to form an ionic compound. Initially, the elements are neutral atoms. In the process of transforming into their ionic compound product, the atoms exchange electrons with one another to make a metal cation (+) and a non-metal anion (–). Generically, we can represent a synthesis reaction as:

A + B → C

And here are some specific reaction examples:

2Na(s) + Cl2(g) → 2NaCl(s)
2H2(g) + O2(g) → 2H2O(l)
CaO(s) + CO2(g) → CaCO3(s)

Decomposition Reactions

A decomposition reaction is one in which a single reactant is broken down into two or more simpler products. Practically speaking, this type of reaction can occur with some sensitive compounds at room temperature but most decomposition reactions require a high temperature. To decompose a compound, heat it up. Other decomposition reactions can be driven by such things as strong ultraviolet light or a catalyst may speed up the normally slow reaction. Some of these reactions are the same as a synthesis reaction written in reverse. Generically, we can give an idea of a decomposition reaction as:

C → A + B

Some specific examples:

CaCO3(s) → CaO(s) + CO2(g)
Mg3N2(s) → 3Mg(s) + N2(g)
2H2O2(aq) → 2H2O(l) + O2(g)

Single Replacement Reactions

In single replacement reactions a neutral pure element is transformed into its ionic form and takes the place of an ion that is part of a compound. The element that began as an ion in turn becomes a neutral pure element. This is a little challenging to describe in words so study the examples and pay attention to whether atoms are alone (neutral) or parts of a compound (ions). The general idea is that one element replaces another within a compound. Here is a general pattern for this type of reaction:

A + BC → AC + B or
A + BC → BA + C

In these reactions it is possible for a metal to replace another metal. It is also possible for a non-metal to replace another non-metal. Normally a metal does not replace a non-metal in a compound, though there is one exception. A metal can replace a hydrogen ion (H+) in an acid, forming hydrogen gas (H2) and an ionic combination of the metal with the anion of the acid. See the examples:

Metal replaces metal:
Cu(s) + 2AgNO3(aq) → Cu(NO3)2(aq) + 2Ag(s)
Non-metal replaces non-metal:
Cl2(g) + 2NaBr(aq) → 2NaCl(aq) + Br2(l)
Metal replaces hydrogen ion:
Fe(s) + 2HCl(aq) → FeCl2(aq) + H2(g)

Double Replacement Reactions

There are two main types of double replacement reactions. Both types involves the exchange of ions between compounds to give two new combinations of ions. In general they looks like this:

AB + CD → AD + CB

In one type of double replacement reaction, an acid reacts with a base to make a salt and water. Acids include the hydrogen ion (H+) in their formula. Bases include the hydroxide ion (OH–) in their formula. These ions combine to make water (H2O). The remaining ions combine to make an ionic compound, also known as a salt. Here are some examples of acid-base reactions:

HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)
H2SO4(aq) + KOH(aq) → 2H2O(l) + K2SO4(aq)

The second type of double replacement reaction involves the exchange of ions between two soluble ionic compounds. Both of the reactants dissolve in water but one of the product ion combinations does not dissolve in water. It is insoluble and because it rains down to the bottom of the beaker it is called the precipitate. You can most readily recognize this type of reaction, called a precipitation reaction, by noticing the phase of matter of the products: one will be dissolved in water (aq) and one will be a solid (s). For example:

CuCl2(aq) + 2NaOH(aq) → 2NaCl(aq) + Cu(OH)2(s)
2Al(NO3)3(aq) +3Na2CO3 → 6NaNO3(aq) + Al2(CO3)3(s)

There is a variation on the theme of exchanging ions which results in one of the products decomposing to make a gas and water. These are a kind of acid-base reaction where the base is a solid compound with a basic anion other than hydroxide. Here are two examples:

CaCO3(s) + 2HNO3(s) → Ca(NO3)2(aq) + CO2(g) + H2O(l)
Na2SO3(s) + 2HCl(s) → 2NaCl(aq) + SO2(g) + H2O(l)

Combustion Reactions

Combustion reactions are reactions in which a compound of hydrogen and carbon (and sometimes also oxygen) combines with oxygen (O2) to make carbon dioxide (CO2) and water (H2O). The reactant other than oxygen can vary but is always a hydrocarbon with molecules including H and C and sometimes O. Combustion reactions are familiar to anyone who has lit a match, sat by a campfire, or given a moment’s thought to how a car is powered or how their home stays warm in the winter. Chemical equations for this type of reaction are more or less instantly recognizable because they always have the same two products (CO2 and H2O) and always have elemental oxygen (O2) as the second reactant. Here are some examples:

CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
2C6H14O(l) + 19O2(g) → 12CO2(g) + 14H2O(g)

Some synthesis reactions release a lot of energy as light and heat. The product of the reaction may even take the form of smoke. But technically these reactions are not considered combustion reactions. Combustion is a word used specifically about the combination of hydrocarbons with oxygen to make carbon dioxide and water.


Objectives

  1. Learn the different types of reactions and recognize them by examining the chemical equation.
  2. Learn how different types of reaction appear in the lab
  3. Learn to connect observations of chemical reactions in the lab with written chemical equations

Pre-lab

Answer each of the following questions in the space provided. When writing chemical equations use what you have learned about chemical nomenclature to write the correct formula. For phases of matter include the following codes after a chemical formula: (s), (l), (g), or (aq). Balance all chemical equations.

  1. Write and balance a chemical equation for a reaction between elemental magnesium metal and aqueous copper(II) nitrate. The products of the reaction are magnesium nitrate and elemental copper metal.
  2. What is the type of reaction for the equation you wrote in the previous question? How did you know?
  3. Write and balance a chemical equation for a reaction between oxygen gas and liquid ethanol (C2H6O). The products of the reaction are carbon dioxide and water vapor.
  4. What is the type of reaction for the equation you wrote in the previous question? How did you know?
  5. Write and balance a chemical equation for a reaction between aqueous sodium phosphate and aqueous iron(II) chloride. The products of the reaction are aqueous sodium chloride and solid iron(II) phosphate.
  6. What is the type of reaction for the equation you wrote in the previous question? How did you know?
  7. Write and balance a chemical equation for a reaction between elemental aluminum metal and liquid elemental bromine (Br2). The product of the reaction is aluminum bromide.
  8. What is the type of reaction for the equation you wrote in the previous question? How did you know?

Materials

Safety

The following list does not cover all possible hazards, just the ones that can be anticipated. Move slowly and carefully in the lab: haste and impatience have caused more than one accident.

Procedure

You may complete the lab stations in any order. Work with your classmates to share access to the materials at each station. Work in pairs and not in larger groups. Clean up after yourself at each station, leaving it ready for the next pair of students to do the experiment.

Chemical Reaction I, Synthesis

  1. At this lab station you will demonstrate for yourself and your lab partner the reaction between magnesium metal and oxygen gas. The magnesium is in the form of thin strips of solid metal that looks like aluminum foil. The oxygen is provided by the atmosphere. Write a balanced chemical equation, including phases of matter, for this reaction. The sole product is solid magnesium oxide.
  2. Make sure the lab station is safe for fire: keep flammable materials away from the lighter or matches and away from the bunsen burner.
  3. Use the lighter or matches to ignite the bunsen burner. If you need help adjusting the flame to a deep blue with a rushing air sound, then ask your teacher. If you have trouble lighting the burner turn off the gas.
  4. Be aware: bright light is produced in the demonstration! Do not stare directly at the burning magnesium. Instead, observe how the light brightens the area around you. Use the tongs to grasp a strip of magnesium. Hold the other end of the strip in the hottest part of the bunsen burner flame. (This is at the top of the inner cone). Hold the burning magnesium over the dish provided for that purpose.
  5. Describe what you observed. Describe the metal, how it ignited, what it was like as it burned, what it produced in the air and on the dish, and how it ended up.
  6. You now have a white material in the dish. Poke it with the end of the tongs and see how it falls apart as a white powder. What is the chemical formula of this powder?
  7. Using your observations and the balanced chemical equation that you wrote, give several pieces of evidence that the reaction you demonstrated is a synthesis reaction.

Chemical Reaction II, Decomposition

  1. At this lab station you will demonstrate for yourself and your lab partner the decomposition of a carbonate compound. The ammonium carbonate is a strong-smelling white powder: do not put your nose in the bottle and breathe in deeply as it will hurt. This substance is corrosive and should be kept away from eyes and mucous membranes. As you heat the ammonium carbonate in a test tube over a bunsen burner three products will be made: water vapor, carbon dioxide gas, and ammonia gas. Write a balanced chemical equation, including phases of matter, for this reaction.
  2. Make sure the lab station is safe for fire: keep flammable materials away from the lighter or matches and away from the bunsen burner.
  3. Use the scoop provided to get a small amount of ammonium carbonate powder. Use no more than about 1 cm of the length of the scoop for your sample. Place your sample in a 1-inch test tube.
  4. Use the lighter or matches to ignite the bunsen burner. If you need help adjusting the flame to a deep blue with a rushing air sound, then ask your teacher. If you have trouble lighting the burner turn off the gas.
  5. Be aware: corrosive ammonia gas is produced in this demonstration! Do not point the test tube at anyone’s face. Use the provided test-tube tongs to grasp the test tube near the opening. Tilt the tube so the opening faces away from you and others nearby and hold the bottom in the hottest part of the bunsen burner flame. (This is at the top of the inner cone). Move the test tube gently in a circle as you heat it.
  6. Observe the cool upper end of the test tube and look for signs of water vapor condensing on the glass on the inside of the test tube.
  7. Use a wet piece of red litmus paper to test for the presence of ammonia gas. Ammonia is a base and turns the color of red litmus to blue. Simply wet the strip in the sink and then stick it to the inside of the opening of the test tube.
  8. Light a wooden splint in the burner flame. Before the ammonium carbonate has completely decomposed into the product gases hold the burning splint inside the test tube. It can be challenging to observe and this may not work. Look for the flame to be extinguished from the burning splint due to the presence of carbon dioxide produce by the decomposition reaction.
  9. Describe what you observed. Describe the powder and its smell. Describe how it behaved in the test tube as you heated it and swirled it around. Describe the tests you did and their results. Were you able to observe all three of the decomposition products?
  10. Using your observations and the balanced chemical equation that you wrote, give several pieces of evidence that the reaction you demonstrated is a decomposition reaction.

Chemical Reaction III, Single Replacement (Acid/Metal)

  1. At this lab station you will demonstrate for yourself and your lab partner the reaction between magnesium metal and hydrochloric acid. The magnesium is in the form of thin strips of solid metal that looks like aluminum foil. The acid is a dilute solution with a concentration of one mole per liter. Write a balanced chemical equation, including phases of matter, for this reaction. The sole products are aqueous magnesium chloride and hydrogen gas.
  2. Make sure the lab station is safe for fire: keep flammable materials away from the lighter or matches.
  3. Take a prepared 3 - 4 cm strip of magnesium metal and roll it up as small as you can. Put it in the bottom of a test tube.
  4. Note: Hydrochloric acid is corrosive and contact with the eyes and mucous membranes should be avoided. If you get some on your skin, just wash with plenty of water. Use a pipet to add about 2 mL of the hydrochloric acid solution to the test tube. A 2-mL volume can be drawn up simply by pressing the bulb of the pipet as flat as you can and drawing up all the liquid you can with it.
  5. As bubbles form in the solution, loosely cover the top of the test tube with your hand to help trap the hydrogen gas.
  6. Pay attention to the temperature of the solution as you hold it against the palm of one hand. You should notice the temperature rising.
  7. Be aware: if you succeed in lighting the hydrogen on fire there will be a loud popping sound. Have your lab partner ignite a wooden splint. When you judge that the reaction has gone on long enough hold the flame over the opening of the test tube. If you do not get it to work and no loud sound occurs, ask your teacher for help. It’s worth it.
  8. Describe what you observed. Describe the metal, how it dissolved, what it was like as it did so, and what happened with the burning splint.
  9. Using your observations and the balanced chemical equation that you wrote, give several pieces of evidence that the reaction you demonstrated is a single replacement reaction.
  10. In this procedure you are meant to observe a loud popping sound as the hydrogen gas ignites, combining with oxygen. Write the balanced chemical equation that represents this reaction. And what type of chemical reaction is it?

Chemical Reaction IV, Single Replacement (Metal/Ionic Compound)

  1. At this lab station you will demonstrate for yourself and your lab partner the reaction between zinc metal and copper(II) sulfate. The zinc is in the form of lumps called ‘mossy zinc’. The copper(II) sulfate is a blue solution with a concentration of one mole per liter. Write a balanced chemical equation, including phases of matter, for this reaction. The sole products are aqueous zinc sulfate and elemental copper powder.
  2. Note: Copper(II) sulfate is toxic by ingestion. Keep away from eyes and mucous membranes. If skin is exposed, wash with plenty of water. Use a graduated cylinder to measure out about 5 mL of the copper(II) sulfate solution. Add it to a test tube secured in a test tube rack.
  3. Take a few small lumps of zinc and add them to the test tube and prepare yourself to wait patiently.
  4. You are looking for the formation of metallic copper but it will take the form of a very fine powder. The powder will form on the surface of the zinc and will intially cling there. It will appear very dark, even black. Later you may observe that it has the reddish-brown color you would expect for copper. Watch for a change in the color of the solution: as the copper changes from its dissolved ionic form to its solid neutral form the solution will lose its blue color.
  5. Describe what you observed. Describe the zinc, how it dissolved, what formed on its surface, and what happened to the color of the solution over time.
  6. The dissolved copper ions have a +2 charge and this is what gives the solution its blue color. What did the zinc atoms have to give to the copper ions in order for the copper ions to lose their charge to become neutral? Explain.
  7. Using your observations and the balanced chemical equation that you wrote, give several pieces of evidence that the reaction you demonstrated is a single replacement reaction.

Chemical Reaction V, Double Replacement (Acid-Gas Generation)

  1. At this lab station you will demonstrate for yourself and your lab partner the reaction between hydrochloric acid and calcium carbonate. The acid is a dilute solution with a concentration of one mole per liter. The calcium carbonate is in the form of a fine white powder. Write a balanced chemical equation, including phases of matter, for this reaction. The products are aqueous calcium chloride, water, and carbon dioxide gas.
  2. Make sure the lab station is safe for fire: keep flammable materials away from the lighter or matches.
  3. Use a metal scoop to get about 1 gram of calcium carbonate powder. This is only a little bit, about a half a teaspoon. Add it to a clean test tube.
  4. Note: Hydrochloric acid is corrosive and contact with the eyes and mucous membranes should be avoided. If you get some on your skin, just wash with plenty of water. Use a pipet to add about 2 mL of the hydrochloric acid solution to the test tube. A 2-mL volume can be drawn up simply by pressing the bulb of the pipet as flat as you can and drawing up all the liquid you can with it.
  5. You are looking for the production of bubbles of gas as the carbon dioxide is generated. Note, also, whether you can feel the temperature of the solution drop as the reaction proceeds.
  6. Have your lab partner ignite a wooden splint. When you judge that the reaction has gone on long enough hold the flame just inside the opening of the test tube. The carbon dioxide is denser than air and will tend to stay in the test tube. If there is enough of it collected then you can detect its presence by the fact that a fire placed into it will go out.
  7. Describe what you observed. Describe the powder, how it dissolved, what it was like as it did so, and what happened with the burning splint.
  8. In this reaction ions are exchanged and initially the products include carbonic acid (H2CO3). This quickly decomposes to form water and carbon dioxide. Write a balanced equation for this decomposition reaction.
  9. Using your observations and the balanced chemical equations that you wrote, give several pieces of evidence that the reaction you demonstrated is a double replacement reaction.

Chemical Reaction VI, Double Replacement (Precipitation)

  1. At this lab station you will demonstrate for yourself and your lab partner the reaction between sodium phosphate and copper(II) sulfate. The sodium phosphate is a colorless solution with a concentration of one half mole per liter. The copper(II) sulfate is a blue solution with a concentration of one mole per liter. Write a balanced chemical equation, including phases of matter, for this reaction. The products are aqueous sodium sulfate and solid copper(II) phosphate powder.
  2. Note: Copper(II) sulfate is toxic by ingestion. Keep away from eyes and mucous membranes. If skin is exposed, wash with plenty of water. Use a pipet to add about 1 mL of the copper(II) sulfate solution to the test tube. A 1-mL volume can be drawn up by filling just the lower part with no liquid in the bulb.
  3. Use a graduated cylinder to add about 10 mL of tap water to the test tube. Swirl to mix. The larger volume makes it easier to observe the results of the reaction.
  4. Note: Sodium phosphate solution is corrosive and contact with the eyes and mucous membranes should be avoided. If you get some on your skin, just wash with plenty of water. Use a pipet to add sodium phosphate solution drop by drop to the copper(II) sulfate solution. Keep adding the clear solution until no more solid products form in the test tube. In all you will add about 2 mL.
  5. You are looking for the appearance of a blue-green powder that forms when the two solutions mix. The particles are very small and will swirl and take a long time to settle to the bottom. After adding the sodium phosphate solution, set the test tube in the rack and wait patiently for the solid to settle to the bottom. If you added enough sodium phosphate the liquid will no longer have a blue color. Take the time to observe this.
  6. Describe what you observed. Describe the solution, how it looked when they mixed together, and the appearance of the liquid in the test tube after the solid settled to the bottom.
  7. Using your observations and the balanced chemical equation that you wrote for the main reaction in this demonstration, give several pieces of evidence that the reaction you demonstrated is a double replacement reaction.
  8. Sodium phosphate is slightly corrosive because the phosphate ion acts as a base in the solution by increasing the concentration of hydroxide ions (OH–). Write a chemical equation for the reaction between aqueous sodium hydroxide and aqueous copper(II) sulfate, which produces solid copper(II) hydroxide and aqueous sodium sulfate. What type of chemical reaction is this secondary reaction?

Chemical Reaction VII, Double Replacement (Acid/Base)

  1. At this lab station you will demonstrate for yourself and your lab partner the reaction between hydrochloric acid and sodium hydroxide. The acid is a dilute solution with a concentration of one mole per liter. The sodium hydroxide is a dilute solution with a concentration of one mole per liter. Write a balanced chemical equation, including phases of matter, for this reaction. The products are aqueous sodium chloride and water.
  2. Note: Hydrochloric acid is corrosive and contact with the eyes and mucous membranes should be avoided. If you get some on your skin, just wash with plenty of water. Use a pipet to add exactly twenty equal-sized drops of the hydrochloric acid solution to the test tube. This is important for the demonstration.
  3. Add one or two drops of the bromothymol blue acid-base indicator solution from the dropper bottle. The solution will turn yellow. For this indicator a solution will be yellow when the pH is 6 or below. The pH of the hydrochloric acid solution starts in this demonstration as pH = 0.
  4. Note: Sodium hydroxide is corrosive and contact with the eyes and mucous membranes should be avoided. If you get some on your skin, just wash with plenty of water. Use a pipet to add the solution to the test tube with the acid and indicator one drop at a time. Swirl the test tube to mix with each drop added. For your information, the sodium hydroxide solution has an initial pH = 14.
  5. You are looking for a permanent color change in the indicator. You will see swirls of color as the solutions mix and different parts of it have different pH values. If you add exactly the same amount of sodium hydroxide as the amount of hydrochloric acid then you will see the color change to green, which shows a neurtal pH (pH = 7). This may be tricky to accomplish but should just be possible.
  6. Keep adding a little more sodium hydroxide after the change to green. The solution will turn blue, which shows that the pH is above 8.
  7. Now close the bottle of sodium hydroxide and get the hydrochloric acid again. Add it by drops to the test tube and swirl to mix. Observe what happens.
  8. Describe what you observed. Describe the solutions, the colors of the indicator and when they shifted, and whether you were able to add exactly the right amount of sodium hydroxide to make the solution green. Also, what happened when you added hydrochloric acid again at the end?
  9. The solutions you mixed in this demonstration had equal concentrations. What if the concentration of the sodium hydroxide had been half as large? If so, how many drops of sodium hydroxide solution would you have needed to add to 20 drop of hydrochloric acid solution to reach the green balance point?
  10. Using your observations and the balanced chemical equation that you wrote, give several pieces of evidence that the reaction you demonstrated is a double replacement reaction.

Chemical Reaction VIII, Combustion

  1. There is no lab station for this demonstration. Rather, combustion can be observed at stations I, II, III, and V. Also, most people have direct experience with candles, matches, lighters, campfires, outdoor grills, and indoor gas stoves. All of these involve the combustion of hydrocarbon fuels by combination with oxygen to make water vapor and carbon dioxide.
  2. Using memory, or a visit to one of the stations involving burning things, describe a flame. Consider color, light, temperature, whether there is smoke or not, and what the fuel looks like as it burns.
  3. Consider that the fuel must be turned into a gas to burn efficiently. With that in mind, draw a picture of a flame and label the parts where you will find fuel and oxygen and the parts where you will find water vapor and carbon dioxide.
  4. Write and balance a chemical equation for the combustion of ethane (C2H6), which is a component in the mixture of gases burned in a bunsen burner.
  5. Write and balance a chemical equation for the combustion of gasoline (one component of the mixture of compounds that make up gasoline is octane: C8H18).

Post-lab

Simply answer all questions and record all observations directly in this lab handout. Turn the paper in to your teacher. This lab assignment must be done by each individual student. You may not copy one another’s work.

Including the pre-lab questions your work will be scored such that every response is worth two points. You may earn partial credit for written responses but blank responses will earn no points.

Last updated: Apr 15, 2025 Home