Experiment 5: Reactions

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EXPERIMENT 5: CHEMICAL REACTIONS

Lansing Community College General Chemistry Laboratory I

LEARNING OBJECTIVES:

Observe changes in chemical properties during a variety of chemical reactions.
Write the molecular, ionic, and net ionic equations for double displacement and single replacement reactions.

INTRODUCTION

There are many kinds of chemical reactions and several ways to classify them. The reaction types include: Combination (Synthesis), Decomposition, Dissociation, Combustion, Single Replacement, and Double Displacement. In this experiment, you will investigate double displacement and single replacement reactions.

Double Replacement Reactions

A double replacement reaction is one in which two compounds, AB and CD, “exchange partners" with each other to form two different compounds. Reactions between two salt solutions and also acid-base neutralizations are double replacement reactions.

General Form: AB + CD → AD + CB

There are three types of chemical equations commonly used to describe a double replacement reaction: molecular, ionic and net ionic equations. There are situations where it is more convenient to use one of these equation types. For example, stoichiometry problems are most easily solved using a molecular equation.

To demonstrate writing equations in all three forms, we will use the reaction between a solution of calcium chloride (CaCl₂) and a solution of sodium carbonate (Na₂CO₃). Analysis of the solid formed would reveal CaCO₃ as the precipitate. After filtering off the solid, the remaining solution, on evaporation, would give a residue (NaCl). The first type of equation we can write for this reaction shows the complete formulas of all reactants and products. We call this equation the molecular equation.

Molecular equation: CaCl₂_(aq) + Na₂CO_3(aq) → CaCO_3(s) + 2NaCl_(aq)

The second type of equation, the ionic equation, gives a more accurate description of the reaction that is taking place. In order to write an ionic equation, we need to be familiar with the way different compounds ionize in aqueous solutions. Compounds that ionize completely or almost completely in aqueous solutions are called strong electrolytes while those that do not ionize or only ionize slightly are known as non-electrolytes and weak electrolytes respectively. When writing ionic equations strong electrolytes should be expressed as ions while weak and non-electrolytes are expressed as molecules. It is also necessary to know the formulas and charges of the ions that exist in solution. Refer to the rules predicting which salts are soluble in water when determining whether a compound is a weak or strong electrolyte, or a non-electrolyte. The ionic equation for the above reaction is as follows. Ions that go through a reaction unchanged and that appear on both sides of the ionic equation are called spectator ions.

Ionic equation: Ca⁺²_(aq) + 2Cl⁻_(aq) + 2Na⁺_(aq) +CO₃^-2_(aq) → CaCO_3(s) + 2Na⁺_(aq) + 2Cl⁻_(aq)

The third and last type of equation, a net ionic equation is more concise and focuses only on ions that undergo a change. Spectator ions (Na⁺ and Cl^-) are omitted.

Net ionic equation: Ca⁺²_(aq) + CO₃^-2_(aq)→ CaCO_3(s)

Another example of double displacement reaction is neutralization reaction. In a neutralization reaction, an acid and a base reacts to form salt of the acid and water. In this type of reaction, the formation of a nonelectrolyte (water) drives the reaction to completion. An example of a neutralization reaction is the reaction of sulfuric acid with sodium hydroxide.

Molecular equation: H₂SO_4(aq) + 2NaOH_(aq) → Na₂SO_4(aq) + 2H₂O_(l)

Ionic equation: 2H⁺_(aq) + SO₄^-2_(aq) + 2Na⁺_(aq) + 2OH⁻(aq) → 2Na⁺ _(aq) + SO₄^-2_(aq)+ 2H₂O_(l)

Net ionic equation: 2H⁺_(aq) + 2OH⁻_(aq) → 2H₂O_(l)

Single Replacement Reactions

In a single replacement reaction, one element takes the place of another element in a compound.

Table of Activity Series
Element	Oxidation
Lithium	Li → Li⁺¹
Potassium	K → K⁺¹
Barium	Ba → Ba⁺²
Calcium	Ca → Ca⁺²
Sodium	Na →Na⁺¹
Magnesium	Mg → Mg⁺²
Aluminum	Al → Al⁺³
Manganese	Mn → Mn⁺²
Zinc	Zn → Zn⁺²
Chromium	Cr → Cr⁺³
Iron	Fe → Fe⁺²
Cobalt	Co → Co⁺²
Nickel	Ni → Ni⁺²
Tin	Sn → Sn⁺²
Lead	Pb → Pb⁺²
Hydrogen	H₂ → 2H⁺
Copper	Cu → Cu⁺²
Silver	Ag → Ag⁺¹
Mercury	Hg → Hg⁺²
Platinum	Pt → Pt⁺²
Gold	Au → Au⁺³

General Form: A + BC → AC + B

In the general form shown above, element A is replacing element B in the compound. The displaced element B is a product in its elemental form. Both metals and nonmetals can be replaced in this manner. A metal will replace another metal (or hydrogen cation), and a nonmetal will replace another nonmetal.

The activity series of metals can be used to predict whether a single replacement reaction involving metal atom (or hydrogen) replacement will occur.

Active metals are effective at displacing cations from compounds. The higher a metal is on the activity series, the more active the metal. The activity series will tell you which metals are sufficiently active to displace hydrogen from an acid. Additionally, a more active metal (higher in the table) can displace a less active metal (lower on the table) from a compound. The metals become ions with the charge shown on the table when they displace a cation.

Example 1: Zn_(s) + 2HCl_(aq) → ZnCl_2(aq) + H_2(g)

In the above example, Zn atoms are displacing H atoms from HCl. Since zinc ions have a +2 charge, the formula for the zinc chloride product is ZnCl₂. The displaced element (hydrogen) will be a product in its elemental form. Since hydrogen is a diatomic gas, 2 H atoms will combine to form a H₂ molecule.

Example 2: Ag_(s)+ H₂SO_4(aq) → No reaction

Silver is not sufficiently high on the activity series to react with acids, so no reaction will occur.

Dissociation of Compounds in Water

We will consider the dissociation of three types of compounds in water: salts, acids, and bases. Water soluble salts, strong acids and strong bases dissociate completely to give ions and therefore are strong electrolytes.

Salts - According to the solubility rules, all chloride salts are water soluble except Pb⁺², Ag⁺ and Hg₂⁺². Therefore, CaCl₂ is a water soluble salt and a strong electrolyte. CaCl₂ will ionize completely in water to give a solution containing calcium cations (Ca⁺²) and chloride anions (Cl⁻).

CaCl_2(aq) → Ca⁺²_(aq) + 2Cl⁻_(aq)

Acids - Seven acids are strong acids: HCl_(aq), HBr_(aq), HI_(aq), HNO₃, H₂SO₄, HClO₃, and HClO₄. Since nitric acid is on this list, it is a strong acid and (like all strong electrolytes) will dissociate completely in water.

HNO_3(aq) → H⁺_(aq)+ NO₃^-_(aq)

Hydrofluoric acid is not one of the seven common strong acids. Therefore, it is a weak acid which only slightly dissociates in water. As weak electrolytes, weak acids ionize in an equilibrium reaction which favors the intact acid molecules with only a small concentration of product ions.

HF_(aq) D H⁺_(aq) + F⁻_(aq)

Bases – Group 1A metal hydroxides, calcium hydroxide, barium hydroxide, and strontium hydroxide are strong bases. Since calcium hydroxide is a strong base it will dissociate completely in water.

Ca(OH)_2(aq) → Ca⁺²_(aq) + 2OH⁻_(aq)

Weak bases only dissociate slightly in water. As weak electrolytes, weak bases ionize in an equilibrium reaction which favors the intact basic compound with only a small concentration of product ions. The general form (where B is a weak base) is:

B + H₂O_(l) D BH⁺_(aq) + OH⁻_(aq)

EXPERIMENTAL PROCEDURE

Part 1 Double Replacement Reactions

You will be mixing two solutions to see if a precipitation reaction occurs. (All solutions are 0.1 M unless stated otherwise.) If precipitation occurs, you will determine the composition of the precipitate by using the strong electrolyte list and solubility rules. Record the observations and complete the molecular, ionic and net ionic equations for each reaction.

Put 20 drops (~1 mL) of an iron (III) nitrate solution into a small test tube. To this solution add 10 drops of 0.50 M sodium hydroxide solution. Complete TABLE 1A.

To another test tube add 20 drops of a lead (II) nitrate solution and 3 drops of 6 M hydrochloric acid. Complete TABLE 1B.1.

Place the test tube containing the lead (II) chloride into the centrifuge. Balance the centrifuge with your neighbor’s or with another test tube filled with water to the same height. Place the tubes opposite to each other. Put the top on the centrifuge. Centrifuge for 2 – 3 minutes. Turn off the centrifuge and wait for it to stop spinning before you open the lid. Decant (pour off) the solution in the lead waste container in the hood. Save this precipitate, which may also contain a very small amount of the supernatant liquid.

To the precipitate in B.1 above add 10 drops of a potassium iodide solution (lead (II) chloride is more soluble than the lead (II) iodide). Complete TABLE 1B.2. Discard all solutions and precipitates containing lead in the labelled lead waste container found in the hood.

Put 20 drops of .050 M Ba(OH)₂ in a test tube. Add a drop of phenolphthalein. Record your observations. Add two drops of 1.0 M HCl. Record your observations and complete TABLE 1C.

Part 2 Single Replacement Reactions

Carry out the following reactions. Record the observations and complete the molecular, ionic and net ionic equations for each reaction.

Use a tweezers to transfer a small piece of magnesium ribbon to a small test tube. Add 10 drops of 1.0 M HCl. Touch and feel the bottom of the test tube. Record your observations and complete TABLE 2A.

Using a spatula add a small amount of iron filings (enough to cover a typed letter “o”) to a clean test tube. To this test tube add 10 drops of 0.50 M CuCl₂ Shake the tube gently. If the blue color disappears, add 10 more drops of 0.50 M CuCl₂. Note the color of the precipitate. Record your observations and complete TABLE 2B.

Part 3 Using Solubility Rules for Precipitate Determination

Note: Have the page with Table 3 in front of you when you are doing this part.

In this part, you will combine pairs of salt solutions and look for precipitates. We will be using a cell well plate (Figure 1).

Figure 1

Place the 24-well micro plate vertically so that you have four columns and six rows. To five wells in column A, add 8 drops of iron (III) nitrate, the first reagent found on the top row of Table 3. Next, add 16 drops of .05 M sodium hydroxide to the first well (upper left) containing iron (III) nitrate. To the next well containing iron (III) nitrate, add 16 drops of sodium sulfate. Continue mixing iron (III) nitrate with the reagents in the left hand column of Table 3. Repeat this procedure for the reagents in the other three columns, making certain that the chemicals are placed in wells that correspond with each solution in Table 3. (All solutions are 0.1 M unless otherwise stated.). You will not be using the sixth row.

If no visible reaction occurred (i.e. no precipitated formed), write “NONE” in the data TABLE 3. If a reaction occurred record the formula of the precipitate formed from the combination of the reagents mixed in that well. Since calcium sulfate has a borderline solubility, wait for about 10 to 15 minutes before recording your results for calcium nitrate + sodium sulfate. For those combination that resulted in a precipitate, write the net ionic equation.

SAFETY NOTE: All lead solutions are toxic. Avoid contact. If you spill any of these solutions on your hands, wash them thoroughly with soap and water.

Clean Up & Waste Disposal

Firmly replace the caps on the reagent bottles immediately after use.
All non-lead containing solutions can be discarded down the drain.
Those that contain lead should be disposed of in the lead waste container found in the hood. Use distilled water to remove the lead precipitate sticking to the test tube. Dispose the washings into the same container.
Wash the well plates with soap and water.
Wipe the bench top with a moist paper towel.

DATA SHEET FOR EXPERIMENT 5 - CHEMICAL REACTIONS

Name ___________________________________ Date: _______________

Part 1: Write a balanced molecular, ionic and net ionic equations for the double displacement reactions that occurred. If a precipitate formed, identify color and texture under “Observations”. Be sure to properly indicate phases (s, l, g, or aq), arrows, charges on ions, and the stoichiometric coefficients (if any) in front of ions to earn full credit.

TABLE 1A: Iron (III) nitrate + Sodium hydroxide (2 pts.)

Molecular
Ionic
Net Ionic
Observations

TABLE 1B.1: Lead (II) nitrate + Hydrochloric acid (2 pts.)

Molecular
Ionic
Net Ionic
Observations

TABLE 1B.2: Precipitate from B.1 + Potassium iodide (2 pts.)

Molecular
Ionic
Net Ionic
Observations

TABLE 1C: Barium hydroxide + Hydrochloric acid (2 pts.)

Molecular
Ionic
Net Ionic
Observations

Part 2: Write the molecular, ionic and net ionic equations for the single replacement reactions that occurred. Be sure to properly indicate phases (s, l, g, or aq), arrows, charges on ions, and the stoichiometric coefficients (if any) in front of ions to earn full credit.

TABLE 2 A: Magnesium + Hydrochloric acid (2 pts.)

Molecular
Ionic
Net Ionic
Observations

TABLE 2 B: Iron + Copper II Chloride (2 pts.)

Molecular
Ionic
Net Ionic
Observations

Part 3: 1. Complete the following table. Record the formulas of the precipitate that forms. If no precipitate forms write “NONE” in the box. (2 pts.)

TABLE 3

A B C D

		Iron(III) nitrate	Copper(II) nitrate	Nickel(II) nitrate	Calcium nitrate
1	Sodium hydroxide
2	Sodium sulfate
3	Sodium chloride
4	Sodium phosphate
5	Sodium carbonate	////OMIT/////// /////////////////////

(2 pts.)For those combinations that resulted in formation of a precipitate in the above table, write the balanced net ionic equations below.

E.g. A1 Fe³⁺_(aq)+ 3OH⁻_(aq)→Fe(OH)_3(s)

POST LABORATORY FOR EXPERIMENT 5 - CHEMICAL REACTIONS

(2 pts.)Balance the following molecular equation and then write the ionic and net ionic equations.

____ Na₂CO_3(aq) + ____ Al(NO₃)_3(aq) → _____ Al₂(CO₃)_3(s) + _____ NaNO_3(aq)

Ionic equation:

Net ionic equation:

(2 pts.)Predict the products and balance the following single replacement reactions. If there is no reaction, write as “No reaction” on the product side.

Na_(s) + H₂SO_4(aq) →

Mg_(s) + HC₂H₃O_2(aq)→

Zn_(s)+ BaCl_2(aq) →

K_(s) + FeCl_2(aq) →

(2 pts.) Write the molecular equation for the reaction, if any that occurs, when the following salt solutions are mixed. (Refer to the solubility table.). If both products are water soluble, then write “No reaction” on the product side.
a) Barium nitrate and sodium chloride
b) Barium nitrate and potassium sulfate