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6: Types of Chemical Reactions (Experiment)

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• Page ID 93998

• Santa Monica College
• To perform and observe the results of a variety of chemical reactions.
• To become familiar with the observable signs of chemical reactions.
• To identify the products formed in chemical reactions and predict when a reaction will occur.
• To write balanced equations for the reactions studied.
• To use the results from the single replacement reactions to devise a partial activity series.

Matter undergoes three kinds of change: physical, chemical, and nuclear. While the composition of a chemical substance is not altered by physical changes (such as freezing and evaporation), chemical changes, or reactions, result in the formation of new substances when bonds are formed and/or broken. Some relatively simple but common types of chemical reactions are illustrated in this experiment. Examples and descriptions of each reaction type appear in the following section.

Chemical reactions are often accompanied by observable changes as they occur. These include:

• Color change.
• Formation of a precipitate—noted as the formation of a cloudy solution, formation of a gel, or
• an obvious solid.
• Evolution of a gas—noted as bubbling in the solution.
• Appearance or disappearance of distinct separation between two or more liquids.
• Evolution of heat—noted as a temperature increase.
• Absorption of heat—noted as a temperature decrease.
• Plating out of one metal on another.
• Decomposition, pitting, or the disappearance of a solid metal

One or more of these changes may occur in the reactions that are performed in this experiment.

Reaction Types

Please note that some reactions may be classified as more than one type of reaction; for example, combination and decomposition reactions that involve elemental substances are also oxidation-reduction reactions.

Combination Reactions (also called Synthesis Reactions) occur when two or more substances, elements or compounds, combine to form one new substance. For example, hydrogen and oxygen gases combine to give water:

\[\ce{2H2 (g) + O2 (g) -> 2H2O(l)}\]

Decomposition Reactions occur when a compound breaks apart to yield two or more new substances. As an example, potassium chlorate decomposes when heated to yield potassium chloride and oxygen gas. Potassium chlorate is one of the ingredients used in match heads. This reaction occurs as the match ignites, causing the match head to be surrounded by an oxygen atmosphere.

\[\ce{2KClO3 (s) -> 2KCl(s) + 3O2 (g)}\]

Displacement Reactions (also called Single Replacement Reactions) involve the displacement of one element in a compound by another element. Several examples of displacement reactions are given below.

• Displacement of one metal by another:

\[\ce{Pb (s) + Cu(NO3)2 (aq) -> Cu (s) + Pb(NO3)2 (aq)}\]

• Displacement of hydrogen gas from an acid by a metal:

\[\ce{Mg(s) + 2HCl(aq) -> MgCl2 (aq) + H2 (g)}\]

• Displacement of hydrogen gas from water by a metal:

\[\ce{2 K (s) + 2 H2O (l) -> 2 KOH (aq) + H2 (g)}\]

• Displacement of one halogen by another:

\[\ce{Cl2 (aq) + NaBr (aq) -> NaCl (aq) + Br2 (aq)}\]

Displacement reactions are also classified as oxidation-reduction reactions. For example, in the first reaction given above, elemental lead is oxidized to lead(II) and copper is reduced from copper(II) to elemental copper. Two electrons are transferred from lead to copper in this process:

\[ \ce{Pb (s) -> Pb^{2+} (aq) + 2 e^-} \quad \quad \text{oxidation of lead}\]

\[ \ce{Cu^{2+} (aq) + 2e^- -> Cu (s)} \quad \quad \text{reduction of copper}\]

The ability of one metal to displace another depends on their relative ease of oxidation—a more active metal (one that is more easily oxidized) displaces a less active metal. In the first reaction above, lead is more active than copper. The relative activities of metals can be tabulated in an activity series, ranking the metals by relative ease of oxidation. A metal that displaces hydrogen gas from acid is more active than hydrogen. A metal that displaces hydrogen gas from acid, but not from water, is less active than one that can displace hydrogen from both acids and water. The ease with which a substance is oxidized is quantified as its standard oxidation potential; you will learn more about this in the second semester of General Chemistry.

Exchange Reactions (also called Double Replacement or Metathesis Reactions) occur when two compounds that form ions in solution react by switching ion partners. Thus, these reactions have the general form:

\[\ce{AB + CD -> AD + CB}\]

One of three conditions must be met for these reactions to occur: (1) the formation of an insoluble ionic compound, observed as a precipitate, (2) the formation of a gas, or (3) the production of water from hydroxide and hydrogen ions (an example of an acid-base neutralization). In each case one of the products results by combining two ions that are removed from the solution by the reaction.

Precipitation Reactions occur when aqueous solutions of two ionic compounds are mixed and the ions combine to make a compound that is insoluble in water (the precipitate). For example, sodium phosphate can be used in an exchange reaction to precipitate calcium ions out of hard water as calcium phosphate, a reaction that is used in some commercial water softeners.

\[\ce{3 CaCl2 (aq) + 2 Na3PO4 (aq) -> Ca3(PO4)2 (s) + 6 NaCl (aq)}\]

The solubility behavior of the ions that you will be using is summarized in the following table:

Gas Forming Reactions typically go to completion because one or more of the products are removed from the reaction vessel via the formation of a gas, which leaves the reaction mixture as bubbles. Gases thus produced include hydrogen sulfide, sulfur dioxide, carbon dioxide and ammonia.

Hydrogen sulfide, \(\ce{H2S}\), is formed by the direct combination of an acid (source of \(\ce{H^{+}}\)) and the sulfide ion:

\[\ce{Na2S(aq) + 2HCl(aq) -> H2S(g) + 2NaCl(aq)}\]

Sulfur dioxide, \(\ce{SO2}\), is formed by the decomposition of sulfurous acid, which is initially formed in a reaction between an acid and the sulfite ion:

\[\ce{Na2SO3 (s) + 2 HCl (aq) -> H2SO3 (aq) + 2 NaCl (aq)}\]

\[\ce{H2SO3 (aq) -> H2O (l) + SO2 (g)}\]

\[\ce{Na2SO3 (s)+2HCl(aq) -> H2O(l) + SO2 (g) + 2NaCl(aq)}\]

Carbon dioxide, \(\ce{CO2}\), is formed by the decomposition of carbonic acid, which is initially formed in a reaction between an acid and the carbonate ion:

\[\ce{Na2CO3 (s) + 2 HCl (aq) -> H2CO3 (aq) + 2 NaCl (aq)}\]

\[\ce{H2CO3 (aq) -> H2O (l) + CO2 (g)}\]

\[\ce{Na2CO3 (s)+2HCl(aq) -> H2O(l) + CO2 (g) + 2NaCl(aq)}\]

Ammonia, \(\ce{NH3}\), is formed from the combination of ammonium and hydroxide ions:

\[\ce{NH4Cl (aq) + NaOH (aq) -> NaCl (aq) + H2O (l) + NH3 (g)}\]

Neutralization Reactions (also called Acid-Base Reactions) involve the transfer of a proton (\(\ce{H^+}\)) from the acid to the base. Water is always a product when the base contains the hydroxide ion (see example below). Some of the gas-forming reactions may also be classified as neutralization reactions.

\[\ce{HCl (aq) + NaOH (aq) -> NaCl (aq) + H2O (l)}\]

Materials and Equipment

Solids: \(\ce{Mg}\), \(\ce{CuSO4*5H2O}\), \(\ce{Ca}\), \(\ce{Cu}\), \(\ce{Zn}\), \(\ce{NaHCO3}\) Solutions: 6 M \(\ce{HCl}\), 6 M \(\ce{NaOH}\), 6 M \(\ce{H2SO4}\), 1 M \(\ce{NH4NO3}\), and 0.1 M solutions of \(\ce{CuSO4}\), \(\ce{ZnSO4}\), \(\ce{AgNO3}\), \(\ce{NaCl}\), \(\ce{Ni(NO3)2}\), \(\ce{Pb(NO3)2}\), and \(\ce{K2CrO4}\)

Equipment: crucible tongs, one large test tube, two small test tubes, ten small test tubes, test tube holder, test tube rack, 100-mL beaker, red litmus paper, Bunsen Burner Supplies for Instructor Demonstrations: \(\ce{CaO}\) solution (prepared in advance by stockroom), sucrose, 18 M \(\ce{H2SO4}\), distilled water, \(\ce{Na}\), means of cutting \(\ce{Na}\) and removing from its storage vessel, two 100-mL beakers, straw, red and blue litmus paper, glass stirring rod

• Do not stare directly at the magnesium when it burns as the light can hurt your eyes.
• Do not touch metals with your hands.

Instructions for Performing each Reaction

Perform each of the following reactions except those that are to be demonstrated by your instructor. You will use 1 mL of solution in many of the reactions; estimate this by drops (typically 12-15 drops from the reagent bottle dispenser) or by measuring 1 mL once in your graduated cylinder and then transferring it to a test tube to determine how far it fills the test tube. Record your observations on the data page as you complete each reaction. Make sure that you observe the results of every reaction even if you didn’t actually mix the chemicals yourself. Then write a balanced equation for each reaction. Be sure to include the states of all compounds in your equations (solid, liquid, aqueous, or gas). If no reaction occurs write the words "no reaction" (or NR) instead of the products in your balanced equation and indicate why your think there was no reaction. Unless otherwise indicated dispose of all waste in the waste container provided. Do not put metal strips in the sink.

Part A: Combination Reactions

• Instructor Demonstration: Pour about 35 mL of a clear saturated solution containing calcium oxide into a 100 mL beaker. Allow the solution to stand for about 15 minutes. Observe. Use a straw to blow bubbles into the solution for a few seconds. Observe the solution again
• Hold a small strip of magnesium metal (used in flashbulbs and fireworks) in your crucible tongs and ignite the metal in the hot portion of a burner flame. Don't forget to note the color and composition of the residue left on the tongs.

Part B: Decomposition Reactions

• Instructor Demonstration: Perform this reaction in the fume hood. Fill a 100 mL beaker about one-third full of granulated sugar (sucrose, \(\ce{C12H22O11}\)). Add about 20 mL of concentrated (18 M) sulfuric acid and stir until mixed well. Continue stirring until the mixture darkens. Observe. Do not touch the reaction products or the beaker with your hands; use a stirring rod to guide the solid product that forms.
• Place a small amount (an amount that will fit on the end of a spatula) of solid copper(II) sulfate pentahydrate in a medium test tube. Use a test tube holder to hold the tube at about a 45 o angle and heat in a burner flame for a few minutes, remembering not to point the tube at anyone in the room. Note any changes in the appearance of the solid and anything else that appears in the test tube. Allow the solid to cool and add a few drops of water. Observe. Dispose of the copper compound in the waste container.

Part C: Displacement Reactions

Use 1 mL of each solution unless otherwise specified. For reactions involving metals, use just one piece of metal. Do not put the metal pieces in the sink. If no discernable initial change is noted, let the reaction mixture stand for at least five to ten minutes before observing again. Not all of the combinations will yield observable reactions. Repeat the reaction if there is any doubt about whether a reaction occurred or not.

• Instructor Demonstration: Cautiously add a small piece of sodium metal to water. Test the resulting solution with red and blue litmus papers (red litmus paper will turn blue in the presence of a base; blue litmus will turn red in the presence of an acid).
• Calcium metal and water (15 mL) in a large test tube
• Zinc metal and water
• Copper metal and 6 M hydrochloric acid
• Zinc metal and 6 M hydrochloric acid
• Zinc metal and 0.1 M copper(II) sulfate
• Copper metal and 0.1 M zinc sulfate
• Copper metal and 0.1 M silver nitrate

Part D: Exchange Reactions

Use 1 mL of each solution unless otherwise specified. Be sure to mix the solutions well.

• 0.1 M silver nitrate and 0.1 M sodium chloride
• 0.1 M nickel(II) nitrate and three drops of 6 M sodium hydroxide
• 0.1 M lead(II) nitrate and 0.1 M potassium chromate
• 1 M ammonium nitrate and 6 M sodium hydroxide. Warm the test tube gently by passing it back and forth through a burner flame. Hold a strip of moistened red litmus paper in the tube without letting it come in contact with the sides of the tube and note any color changes to the paper. Remove the tube from the flame and quickly and cautiously note the smell.
• Place 5 mL of 6 M hydrochloric acid in a 100 mL beaker. Carefully add several spatulas of solid sodium bicarbonate. Observe.
• Combine about 5 mL each of 6 M sodium hydroxide and 6 M sulfuric acid in a large test tube. Mix with a stirring rod. Cautiously feel the outside of the test tube. If you cannot detect anything, make sure that you used the correct concentrations of acid and base. Be very careful with the concentrated sulfuric acid; it is very caustic and can dissolve skin and clothing.

Pre-laboratory Assignment: Types of Reactions

• Many of the reactions use 1 mL of solution. How can you estimate this volume?
• Do you need to dry the test tubes before using them for the reactions in this experiment? Why or why not?
• The following reactions are performed, and the results are shown below. Use these results to determine the relative activities of the two elements involved in each reaction next to that reaction. Then place the elements gold, hydrogen, zinc and tin in an activity series in order of decreasing activity.

\[\ce{Sn(s) + HCl(aq) -> H2(g) + SnCl2(aq)} \quad \quad \quad \text{______ >______}\]

\[\ce{Au (s) + Sn(NO3)2 (aq) ->} \text{NR} \quad \quad \quad \quad \text{______ >______}\]

\[\ce{Au(s) + HCl ->} \text{NR} \quad \quad\quad \quad \text{______ >______}\]

\[\ce{Zn (s) + Sn(NO3)2 (aq) -> Zn(NO3)2 (aq) + Sn (s)} \quad \text{______ >______}\]

From most active (most easily oxidized) to least active:

______>______>______>______

Now use the above results to write products for the reactions below. Write NR if no reaction is expected.

\[\ce{Sn (s) + Zn(NO3)2 (aq) ->}\]

\[\ce{Zn (s) + Au(NO3)3 (aq) ->}\]

• Suppose that each of the following pairs of aqueous solutions is combined. For those where a reaction is expected, write a balanced formula equation, with state labels, for the reaction that occurs. If no reaction is expected, indicate this and explain why no reaction is expected.
• Barium chloride + potassium sulfate
• Aluminum nitrate + sodium chloride
• Sodium hydroxide + phosphoric acid

Lab Report: Types of Reactions

Record your observations on these data pages as you perform each reaction. Write a balanced formula equation with state labels for each reaction. If no reaction occurs, follow the instructions in the Procedure.

• Calcium oxide and carbon dioxide (demonstration)
• Magnesium and oxygen
• Sucrose and sulfuric acid catalyst (demonstration)
• Thermal decomposition of copper(II) sulfate pentahydrate, \(\ce{CuSO4*5H2O}\)

In addition to providing observations and an equation for each reaction, use your results to determine the relative activities of the two elements involved in each reaction.

• Sodium and water (demonstration)
• Calcium and water
• Zinc and water
• Copper and hydrochloric acid
• Zinc and hydrochloric acid
• Zinc and copper(II) sulfate
• Copper and zinc sulfate
• Copper and silver nitrate

Arrange copper, silver, calcium, zinc, and hydrogen in an activity series from most active to least active on the basis of the results from the displacement reactions that you performed. Recall that a more active metal displaces a less active metal, a more active metal to is needed to displace hydrogen from water than to displace it from an acid, and that a metal that displaces hydrogen from acid is ranked as more active than hydrogen.

________>________>________>________>________

most active (most easily oxidized) to least active

• Silver nitrate and sodium chloride
• Nickel(II) nitrate and sodium hydroxide
• Lead(II) nitrate and potassium chromate
• Ammonium nitrate and sodium hydroxide
• Hydrochloric acid and sodium bicarbonate
• Sulfuric acid and sodium hydroxide

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Assignment Chemical Reactions Lab Report