Unit 4 Study Guide

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In this unit you will learn how to use the coefficients in a balanced equation to calculate: (1) the amount of reactants required for a reaction and (2) the theoretical yield of products expected from a reaction. You will also learn: (3) how to analyze a situation in which one of the reactants is in excess and (4) the concepts of (molar) concentration and making solutions of desired volume and concentration. The Concept of Stoichiometry

Kitchen Stoichiometry (an analogy from your everyday experience) A balanced equation is like a RECIPE

1. I want to make 10 cups of trail mix. How many cups of almonds must I buy?

(3.33 cups) 2. I want to make 3 cups of trail mix. The health food store sells raisins by mass

(300g = 1 cup). How many grams of raisins must I buy?

(150 g = 0.50 cups) 3. I only have 1.5 cups of almonds. (I have plenty of raisins and granola in the house or what we

would call an excess of raisins and granola) a. How many cups of trail mix can I make?

(4.5 cups) b. How many cups of raisins will I need for this mix?

(0.75 cups)

Unit 4: Stoichiometry and Solution Concentration

The Recipe: 2C Almonds + 1C Raisins + 3C Granola 6C Trail Mix

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Stoichiometry and Chemical Reactions You have learned to balance equations by inspection using a guided trial and error technique to arrive at a proper set of coefficients. For example:

N2 + 3 H2 2 NH3

How can this set of coefficients be interpreted?

On the particle level: N2 + 3 H2 2 NH3

1 molecule of N2 reacts with 3 molecules of H2 to produce 2 molecules of NH3 That's valid, but experimental chemistry can't be done on the particle level. Remembering that the mole is a way of describing a macroscopic sample of particles, it is more useful to interpret the coefficients on the macroscopic level: N2 + 3 H2 2 NH3

1 mole of N2 reacts with 3 moles of H2 to produce 2 moles of NH3

IN-CLASS EXAMPLES (using the equation above)

One reactant can be related to another reactant. Example 1. How many moles of nitrogen gas will be needed to completely react with 6.0 moles of hydrogen gas?

(2.0 moles) A reactant and a product can be related. Example 2. How many moles of hydrogen gas will be needed to produce 3.0 moles of ammonia gas?

(4.5 moles) Example 3. How many moles of ammonia gas can be produced from 0.867 moles of nitrogen gas?

(1.73 moles) OR one product could be related to another product using the same concept.

THERE IS ONE MAIN CONCEPT: WHEN QUANTITATIVELY RELATING TWO SUBSTANCES IN A REACTION, THE COEFFICIENTS OF THE SUBSTANCES WILL BE IN THE SAME PROPORTION AS THE MOLE AMOUNTS OF THE SUBSTANCES.

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Stoichiometry Problems Involving Mass In more conventional stoichiometry problems, the given information and/or the requested answer will be in the form of the mass of the substances. You will have to get the MM involved. There are several ways to solve these problems and different teachers may prefer different set ups or strategies.

IN-CLASS EXAMPLE

Example 4. For the reaction: N2 + 3H2 2NH3

What mass of hydrogen gas is needed to produce 10.2 g of ammonia?

[1.8 g (or 1.82 g if MM H2 = 2.02 g/mol)] Stoichiometry Problems

4. Silver reacts with nitric acid (HNO3) to form nitrogen monoxide, silver nitrate and water.

3 Ag + 4 HNO3 → NO + 3 AgNO3 + 2 H2O

a. How many moles of nitric acid are required to react with 5.00 grams of silver?

(0.0618 mol) b. How many grams of silver nitrate will be formed?

(7.87 g)

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5. Hydrazine (N2H4) and hydrogen peroxide (H2O2) react exothermically to produce nitrogen and water. This combination is sometimes used as rocket fuel.

N2H4 + 2 H2O2 → N2 + 4 H2O

How many grams of hydrogen peroxide are needed to react with 100.0 grams of hydrazine?

(212.3 g)

6. When zinc is reacted with phosphoric acid (H3PO4), zinc phosphate and hydrogen are produced. a. Write the balanced molecular equation for this reaction.

___________________________________________________________________________ (answer to balanced equation on p.39)

b. How many moles of phosphoric acid are required to completely react with 17.5 grams of zinc?

(0.179 moles) c. What mass of zinc phosphate will be produced when 17.5 grams of zinc react?

(34.4 g)

7.a. Write the balanced equation for the decomposition of water into its elements.

____________________________________________________________________________ (answer to balanced equation on p.39)

b. How many moles of oxygen are produced by the decomposition (electrolysis) of 108.0g of water?

(3.00 moles)

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8.a. Write a balanced synthesis equation for the production of rust (Fe2O3).

____________________________________________________________________________ (answer to balanced equation on p.39)

b. What mass of iron will produce 11.2 grams of rust?

(7.83 g) c. What mass of oxygen will combine with the iron in part (a)?

(3.37 g) 9. For the reaction: 2 Na(s) + Zn(NO3)2 (aq) → 2 NaNO3(aq) + Zn(s)

a. How many grams of sodium are needed to react completely with 63.13g of zinc nitrate?

(15.3 g) b. In another experiment using this reaction 15.88g of Zn are produced, what mass of sodium was reacted?

(11.2 g)

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Theoretical and Actual Yield Most of this is review. You have used these concepts before, but we can now define theoretical yield in terms of stoichiometry. From your work with stoichiometry problems, you have learned to theoretically predict the mass of products from the given mass of a reactant. However, you have also learned from your lab experience, that often the amount of product obtained in an actual experiment is more or less than the amount that should have been obtained. This happens because the design and/or the techniques used in the experiment are not perfect. The theoretical yield of a given product is the maximum amount of product that should be obtained in a given experiment based on the amount or reactants used. In other words, the theoretical yield is the calculated yield predicted by stoichiometry. The actual yield is the amount of product actually obtained in a given experiment. Normally, it is obtained by isolating the product of the reaction and weighing it. The percent yield is a measure of the efficiency of the reaction and is defined as:

actual experimental yield% yield = (100)

theoretical yield

When an “actual yields” is greater than the “theoretical yield” it usually implies an impure product. If the actual yield is less than the theoretical yield, it is usually because of incomplete reaction, side reactions or loss of product in one of the steps in the experiment.

Problems on % Yield

10. In an experiment involving Zn and HCl, 140.15 g of ZnCl2 was actually formed. The theoretical yield was predicted to be 143.0 g. Calculate the percent yield.

(98.0%) 11. Copper reacts with chlorine gas to synthesize copper(II) chloride. a. Write a balanced equation for this reaction.

________________________________________________________________________ (answer to balanced equation on p.39)

b. If 12.5 g of copper is reacted, calculate the theoretical yield of copper(II) chloride.

c. If 25.4 g of copper(II) chloride is produced in the experiment, calculate the percent yield.

(b. 26.5 g c. 95.8%)

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Limiting Reactant Stoichiometry

Back to the Kitchen Stoichiometry analogy:

12. I only have 3C of almonds and 4C of granola in the house and plenty of raisins.

a. Which of these two ingredients (almonds or granola) is the “limiting reactant” (1.e. the ingredient that will be entirely used up)?

(granola) b. How much of the "excess" ingredient (almonds or granola) will remain after the entire limiting

ingredient” is used?

(0.33 cups of almonds will be left over) c. How many cups of trail mix can I make?

(8 cups)

The Recipe: 2C Almonds + 1C Raisins + 3C Granola → 6C Trail Mix

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Limiting Reactant Stoichiometry and Chemical Reactions

If two reactants are NOT combined in the proper mole ratio (Mother Nature's "recipe"), the reaction will stop when one of the reactants is completely consumed (limiting reactant = LR). The other reactant will be left over (reactant in excess = XS). The final amount of product will be determined by the initial amount of limiting reactant.

There are a few different strategies for setting up and solving these problems. Each teacher will have a preference.

The following examples will allow you to work on these strategies: The first one has given information in moles and asks for answers in moles. The second one is a more conventional problem with given information and requested answers as masses.

IN-CLASS EXAMPLES

For the reaction: N2 + 3 H2 2 NH3

Example 5. If 2.0 moles of N2 are combined with 7.0 moles of H2: a. Which reactant is the limiting reactant?

(N2 is the limiting reactant)

b. How many moles of the excess reactant will be left over?

(1.0 mol of H2 will be left over.)

c. How many moles of NH3 will be produced?

(4.0 mol)

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N2 + 3H2 2NH3 Example 6. If 8.00 g of hydrogen gas and 70.0 g of nitrogen are mixed: a. Which reactant is the limiting reactant?

(H2 is the limiting reactant)

b. What mass of the excess reactant will be left over?

(32.8 g of N2 will be left over.)

c. What mass of NH3 will be produced?

(45.4 g)

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Limiting Reactant Stoichiometry Problems

13. Nitrogen monoxide gas reacts with oxygen to form nitrogen dioxide gas.

2 NO + O2 2 NO2

a. If 0.36 mol of nitrogen monoxide and 0.25 mol of oxygen are combined, which reactant is the limiting reactant?

(NO) b. How many moles of nitrogen dioxide will be produced?

(0.36 mol) c. How many moles of the excess reactant will be left over?

(0.07 mol O2)

d. If the amount of nitrogen monoxide were increased, would the yield of nitrogen dioxide increase, decrease or remain the same? Why?

______________________________________________________________________________ ______________________________________________________________________________

(answer on p. 39) e. If the amount of oxygen were increased, would the yield of nitrogen dioxide increase, decrease or remain the same? Why?

______________________________________________________________________________ ______________________________________________________________________________

(answer on p. 39)

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14. During World War I, the substance phosphine (PH3) was used as a poisonous gas against the allied troops in their trenches (phosphine is heavier than air). Phosphine may be produced by the reaction:

Na3P(s) + 3 H2O(l) → PH3(g) + 3 NaOH(aq)

If 165 g of sodium phosphide is reacted in 255 g of water,

a. Which reactant is the limiting reactant?

(Na3P) b. What mass of each reactant will remain after the reaction is complete?

(Na3P – None left over) (H2O – 166 g left over)

c. How many grams of phosphine (PH3) can be produced by this mixture?

(56.1 g)

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15. Hydrogen peroxide is used as a cleaning agent in the treatment of cuts and abrasions for several reasons. It is an oxidizing agent that can directly kill many microorganisms; it decomposes upon contact with blood, releasing elemental oxygen gas (which inhibits the growth of anaerobic microorganisms); and it foams upon contact with blood, which provides a cleansing action. In the laboratory, small quantities of hydrogen peroxide can be prepared by the action of an acid on an alkaline earth metal peroxide, such as barium peroxide.

BaO2(s) + 2 HCl(aq) → H2O2(aq) + BaCl2(aq)

A student mixed 1.50 g of barium peroxide with a sample of a hydrochloric acid solution that contained 0.816 g of HCl.

a. Which is the limiting reactant?

(BaO2) b. What mass of the excess reactant remains after a reaction is complete?

(0.17 g HCl) c. What mass of hydrogen peroxide can be produced by this combination?

(0.301 g)

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16. In the thermit reaction: iron(III) oxide reacts with aluminum to produce aluminum oxide and iron. What is the theoretical yield of iron from a mixture of 15.0 g aluminum and 25.0 g of iron(III) oxide?

Fe2O3 + 2 Al Al2O3 + 2 Fe

(17.5 g Fe)

17. When 10.0 g of hydrogen is combined with 10.0 g of chlorine and the mixture is exposed to a flash of very bright light, hydrogen chloride is produced with a big “bang”. The amount of hydrogen chloride recovered from this process is 9.8 g.

H2(g) + Cl2(g) 2 HCl(g)

What is the % yield of hydrogen chloride?

(95.1 %)

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Volumetric Flask

Solution Concentration: Molarity

Concentration describes how much solute is in a given amount of solution. There are many ways to describe concentration. For example, 2% milk contains 2 grams of fat for every 100 grams of milk. 70% isopropyl alcohol contains 70 milliliters of alcohol for every 100 milliliters of solution (the rest is water).

Many of the reagents we use in chemistry are dissolved in water. The concentration of a solution can be expressed in a variety of ways. For example, suppose we wanted to prepare a sodium chloride solution. We could mass a certain amount of NaCl, say 20.0 grams, and dissolve it in enough water to make a liter of solution. The concentration of that solution would be 20.0g NaCl/liter of solution. The problem with this concentration unit (grams/liter) is that it does not directly convey the number of particles dissolved in solution and chemical reactions fundamentally occur between particles. Chemists express solution concentration in terms of the molarity (abbreviated M) which directly describes the number of particles dissolved in a given volume of solution.

moles of solute Molarity =

liters of solution

n M =

V

A 1.0 M (1.0 "molar") solution contains 1 mole of solute per liter of solution, a 2.5 M (2.5 molar) solution contains 2.5 moles of solute per liter of solution, etc.

IN-CLASS EXAMPLES

Example 7

Calculate the concentration of a solution made by putting 0.50 moles of sodium chloride in a flask and adding water to the flask until the volume of the solution is 250 mL.

(2.0 M NaCl) Example 8

What is the concentration of a solution prepared by mixing 15.0 grams of NaCl with enough water to make a 4.00 liter solution?

(6.41 x 10-2 M NaCl)

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The relationship shown below (in three different forms) can be used to solve for moles of a solute in a given volume of solution of known concentration. It's not new; it's simply a rearrangement of the definition of molarity:

Example 9 Suppose you have 0.500 L of a 2.20 M solution of sodium chloride. a. If you separate the water from the salt by distillation, how many moles of NaCl would remain in the container?

(1.10 moles) b. What mass of NaCl would remain?

(64.4g) Example 10

Give the directions for making 3.00 liters of a 2.40M solution of NaCl. In other words, calculate the mass of solute needed to prepare of the solution and describe how you would determine the amount of water required.

(421 g of NaCl and enough water so that the volume of the solution is 3.00 liters) Problems on Molarity and Preparing Solutions

18. What is the concentration of each of these solutions? a. 0.175 moles CoCl2 in 350.0 mL of solution.

(0.500 M) b. 8.00 g NaOH in 2.00 liter of solution.

(0.100 M) c. 3.23 g ZnSO4 in 50.0 mL of solution.

(0.400 M)

or or(in liters) moles

moles = x Liters moles = Molarity x volume n = MVLiter

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19. How many moles of solute are contained in each of the following sucrose solutions? How many grams of sucrose (C12H22O11) are in each solution?

a. 1.00 L of a 2.53 M solution.

(2.53 mol = 865 g) b. 500.0 mL of a 0.132 M solution.

(0.0660 mol = 22.6 g) c. 4.0 mL of a 1.30 M solution.

(0.0052 mol = 1.8 g)

20. Calculate the mass of solute needed for the preparation of the following solutions and describe how you would determine the amount of water required.

a. 2.00 liters of a 1.50 M KOH solution.

(168 g of KOH; add water until volume of solution is 2.00 L) b. 1.50 L of a 0.0135 M ZnSO4 solution.

(3.26 g ZnSO4 and enough water to make 1.50 L of solution) c. 35 mL of a 1.3 M Fe(NO3)2 solution.

(8.2 g in 35 mL solution)

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Dilution of Solutions A dilute solution (lower concentration) can be made by adding water to a more concentrated solution or "stock" solution. The major concepts to remember when water is added to a sample of a concentrated solution to produce a dilute solution are:

1) The sample of concentrated solution and the sample of dilute solution contain the same number of moles of solute:

solute c solute d

c c d d

n = n

M V = M V

2) The volume of the dilute solution is equal to the volume of concentrated solution plus the volume of water added.

2d c H OV = V + V

Problems on Dilution of Solutions

21. How would you prepare 1.00 L of 0.646 M HCl from a 2.00 M HCl solution? (i.e. What volume of the concentrated solution is needed, and what volume of water is needed?)

(0.323 L of 2.00 M HCl and 0.677 L of water)

22. What is the concentration of a relatively dilute solution of KNO3 that is prepared by adding 175 mL of water to 25.0 mL of 0.866 M KNO3?

(0.108 M KNO3)

23. You have 505 ml of a 0.125 M NaOH solution. You want to dilute it to exactly 0.100 M NaOH. a. What will the volume of the dilute solution be?

[0.631 L (or 631 mL)] b. How much water should you add to the concentrated solution?

[0.126 L (or 126 mL)]

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Solution Stoichiometry

Many chemical reactions occur between substances dissolved in water (aqueous solutions). You have seen solids precipitate when certain aqueous solutions are mixed. You have added aluminum metal to an aqueous solution of copper ions to produce copper metal and aqueous aluminum ions. The reactants (and products) of acid-base double replacement reactions and redox reactions are often in aqueous solution.

Now that you are familiar with molarity and stoichiometry problems, you are ready to combine the two concepts and learn how to do stoichiometry calculations when the reactant(s) are dissolved in water. This is not entirely new; it's a variation of the concept we've been using.

The definition of molarity n

MV

can be used to calculate moles, molarity, or volume.

IN-CLASS EXAMPLES Example 11

When liquid bromine is mixed with a solution of potassium iodine, aqueous iodine and potassium bromine are produced. The equation for this reaction is:

Br2(l) + 2 KI(aq) I2(aq) + 2 KBr(aq)

What mass of bromine will react completely with 25.0 mL of 1.50 M KI?

(3.00 grams) Example 12

a. Balance molecular equation for the reaction between hydrochloric acid and barium hydroxide.

HCl(aq) + Ba(OH)2(aq) BaCl2(aq) + H2O(l)

b. 24.6 mL of 0.600 M hydrochloric acid are needed to completely neutralize 20.0 mL of a barium hydroxide solution of unknown concentration. Calculate the concentration of the barium hydroxide solution.

(a. 2 HCl(aq) + Ba(OH)2(aq) BaCl2(aq) + 2 H2O(l) / b. 0.369 M)

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Solution Stoichiometry Problems

Use the balanced equation below to answer the next two problems (24-25).

Pb(NO3)2(aq) + 2 NaCl(aq) → 2 NaNO3(aq) + PbCl2(s)

24. What mass of solid lead(II) chloride can be produced when 120.0 mL of 0.200 M lead(II) nitrate is mixed with an excess of sodium chloride?

(6.68 g) 25. What volume of 0.600 M sodium chloride is needed to produce 10.3 grams of PbCl2? (assume an excess of lead nitrate).

[0.123 L (=123 mL)]

26. a. Complete and balance the following equation. NaOH(aq) + H2CO3(aq) _________________________________

(answer to balanced equation on p. 39)

b. 17.3 mL of 1.30 M H2CO3 are needed to completely neutralize 25.0 mL of a sodium hydroxide solution of unknown concentration. Calculate the concentration of the sodium hydroxide solution.

(1.80 M)

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27. In an acidic solution nitrate ions and chloride ions react according to the following skeletal redox reaction:

NO3- + Cl- → NO + Cl2

a. Balance the two half reactions and write the balanced equation on the line below.

(answer to balanced equation on p.39)

b. 11.6 mL of a solution that is 0.500 M in nitrate ions are needed to completely react with 10.0 mL of a solution containing an unknown concentration of chloride ions. Calculate the concentration of the chloride ions in the original solution.

(1.74 M)

28. a. Use the balanced equation below to answer the next problem.

Hg(NO3)2(aq) + 2 KI(aq) → 2 KNO3(aq) + HgI2(s)

If 500.0 mL of 0.200M mercury(II) nitrate is mixed with 250.0 mL of 0.600M potassium iodide, which reactant will be the limiting reactant?

(potassium iodide)

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(DEMONSTRATION)

BACKGROUND One of the most appetizing smells in the world might be that of freshly baked cakes and muffins. In order to make such tasty morsels light and fluffy, the substance known as sodium bicarbonate (NaHCO3) must be included in the recipe. Sodium bicarbonate (baking soda, as it is more commonly named) is the substance that causes many baked goods to rise. This rising action occurs because when strongly heated in the oven, Baking soda decomposes producing one (or more) gases. The gas becomes trapped and produces a fluffy mixture.

A problem many beginning chemistry students face is the ability to predict the products for a particular reaction. The decomposition of sodium bicarbonate presents an interesting dilemma. This is interesting, because it seems to have five plausible, but different products as shown below.

A. NaHCO3(s) → NaOH(s) + CO2(g)

B. NaHCO3(s) → Na(s) + C(s) + H2(g) + O2(g)

C. NaHCO3(s) → Na2CO3(s) + CO2(g) + H2O(g)

D. NaHCO3(s) → Na2O(s) + CO2(g) + H2O(g)

E. NaHCO3(s) → NaH(s) + CO2(g) +O2(g) Your task in this exercise is to determine which of the above equations predicts the correct

products. PRELAB ASSIGNMENT Write a Purpose on page 23. Balance all of the above equations before coming to class, and copy them into the Predictions / Results Table on page 23. METHOD First, we will determine whether the gas(es) produced by the decomposition of NaHCO3 is CO2 or H2 and O2. A reliable test for the presence of carbon dioxide is the “limewater” test. Limewater is a solution of calcium hydroxide. When carbon dioxide is bubbled through a calcium hydroxide solution, a precipitate of calcium carbonate is formed.

Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)

When hydrogen and/or oxygen gas is bubbled through a solution of calcium hydroxide, there is no reaction.

In the second part of this exercise, we will decompose a known mass of the NaHCO3 and determine the actual mass of solid product that is formed.

Lab 4-1: The Bicarbonate Dilemma

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In the analysis, you will use stoichiometry and each of the possible equations to predict the mass of solid product that would be formed by each reaction. By comparing the actual mass of solid product to the predicted masses you can determine the correct reaction. DEMONSTRATION PROCEDURE 1. Put some sodium bicarbonate in a test tube. Stopper the test tube with a rubber stopper into which a bent glass tube has been inserted. Assemble the apparatus as pictured on the right. Pour about 5-10mL of a saturates solution of Ca(OH)2 (limewater) into a second test tube and put it aside for the moment. Begin heating your NaHCO3, then insert the bent glass tube into the limewater solution. Record the observations and discard the sodium bicarbonate and the limewater in the sink with plenty of water. 2. Mass a clean, dry crucible, and then add about 3-4 grams of NaHCO3. Record the exact mass. Assemble a tripod, clay triangle and your crucible and heat the NaHCO3 strongly for about three minutes. Allow the crucible to cool (cool enough so that it can be transferred to the balance), and mass. Repeat the heating and cooling processes until the mass is constant ( 0.02g). 3. CLEAN UP…NaHCO3 and residual solids can all go in the sink. Do not wash the crucible. Wipe it clean with a paper towel. ANALYSIS AND CONCLUSIONS:

Follow the instructions on the worksheet to perform the following calculations and answer the following questions. 1. Check that you have written a balanced equation for each of the reactions on p. 23.

2. Identify the gas produced by the decomposition of NaHCO3. Does identification of the gas eliminate any of the possible equation(s) for this reaction?

3. Calculate the moles of sodium bicarbonate used in this experiment.

4. Calculate the predicted / theoretical mass of the solid product(s) which should have been produced in each of the reactions not eliminated in question 2. SHOW ALL OF YOUR WORK.

5. Calculate the mass of solid product(s) actually formed in the experiment.

6. Compare the predicted / theoretical mass of solid product(s) in each of the possible reactions to the actual mass of the solid product(s), and determine which of the reactions seems to be the most probable.

7. Compare your experimental mass of the product with the theoretical mass of the product determined by stoichiometry by calculating the percent yield.

8. Fill in the tables on p. 23.

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NAME _____________________________________

DATE____________________

LAB 4-1: BICARBONATE DILEMMA PURPOSE: ___________________________________________________________________ _____________________________________________________________________________ OBSERVATIONS:

Limewater Test:

DATA:

Object / Substance mass (grams)

crucible

NaHCO3

crucible & solid product(s) after 1st heating

crucible & solid product(s) after 2nd heating

PREDICTIONS / RESULTS:

Possible reaction

Balanced Equation Theoretical Yield

predicted mass of solid product(s) (grams)

A

B

C

D

E

Actual Yield actual mass of solid product (g)

equation for most probable rxn

% yield

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ANALYSIS 1. Done before class 2. Gas produced ________________ Possible reaction(s) eliminated __________________ 3. Calculate the moles of NaHCO3 used in the experiment. 4. For each reaction that is not eliminated by the limewater test, write the balanced equation, and calculate the predicted mass of the solid product(s) in the allotted space. RXN A. (Question 4) Write the balanced equation for possible reaction (A) on the line below, then calculate the predicted mass of the solid product(s) formed in reaction (A). ______________________________________________________________________________

RXN B. (Question 4) Write the balanced equation for possible reaction (B) on the line below, then calculate the predicted mass of the solid product(s) formed in reaction (B). ______________________________________________________________________________

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RXN C. (Question 4) Write the balanced equation for possible reaction (C) on the line below, then calculate the predicted mass of the solid product(s) formed in reaction (C). ______________________________________________________________________________

RXN D. (Question 4) Write the balanced equation for possible reaction (D) on the line below, then calculate the predicted mass of the solid product(s) formed in reaction (D). ______________________________________________________________________________

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RXN E. (Question 4) Write the balanced equation for possible reaction (E) on the line below, then calculate the predicted mass of the solid product(s) formed in reaction (E). ______________________________________________________________________________

5. Calculate the mass of solid product actually formed. 6. Write the balanced equation for the most probable reaction in the results table. 7. Calculate the percent yield of solid products in this experiment.

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(INDIVIDUAL)

METHOD

In this experiment you will be observing the reaction between calcium chloride and sodium carbonate. The equation for this reaction is:

1 CaCl2(aq) + 1 Na2CO3(aq) → 1 CaCO3(s) + 2NaCl(aq)

Since both reactants are soluble in water, the reaction can be accomplished by mixing solutions of these compounds.

The products of the reaction are calcium carbonate and sodium chloride. These two products can be easily separated from each other because one is soluble in water and the other is not. Calcium carbonate is insoluble in water so it will form a precipitate which can be collected by filtration and evaporating all the water. The sodium chloride can then be massed.

In addition to observing this precipitation reaction, you will carry out a number of reaction calculations. You will predict the theoretical yields of both products. You will then compare the actual yields of both products to their theoretical yields to obtain percent yields. PRELAB ASSIGNMENT

Notice that in the balanced equation for this reaction the mole ratio of CaCl2 to Na2CO3 in this reaction is 1:1. In order to ensure that neither reactant is present in excess, you will want to mass equal moles of CaCl2∙2H2O and Na2CO3∙H2O. For this pre-lab assignment, you will calculate the exact mass (in grams) of each reactant to be used in this experiment by following the steps below:

1) Pick a mass between 1.00 g and 1.50 g. This is the mass of Na2CO3 ∙ H2O you will use in this experiment.

2) Convert the mass of Na2CO3 ∙ H2O from step 1 to moles of Na2CO3 ∙ H2O. 3) Based on the moles of Na2CO3 ∙ H2O from step 2, calculate the exact number of moles of

CaCl2 ∙ 2 H2O you will need to use in this experiment for a stoichiometric reaction. 4) Convert the moles of CaCl2 ∙ 2 H2O from step 3 to grams of CaCl2 ∙ 2 H2O. 5) Fill in the table provided below.

Show calculations below (report your answers to 3 significant digits): ** Notice that both reactants are hydrates. The integrated water must be taken into account when calculating the molar mass of each reactant.

Reactant Mass (g)

Moles (moles)

Na2CO3 ∙ H2O

CaCl2 ∙ 2 H2O

Lab 4-2: Synthesis and Separation of Calcium Carbonate and Sodium Chloride

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PROCEDURE Either before or after massing your reactants, heat the evaporating dish and watch glass for 3 minutes to make sure they are dry when you mass them prior to performing step 7.

1. Mass equal mole amounts of the two reactants into separate beakers. 2. Dissolve each compound in about 10 mL distilled water. (Be patient; the sodium carbonate

dissolves slowly. Gentle heating speeds up the dissolving.) 3. Combine these solutions in the smaller beaker. 4. Rinse the large beaker with your wash bottle and add the rinsings to the small beaker. 5. Swirl the contents for about one minute and write your observations. 6. Obtain a piece of filter paper in order to carry out the filtration described below.

Filtration- Separating a Solid from a Solution

You have done this before, but this will serve as a reminder. You need to fold the filter paper twice, put our name on it and mass it. Then open the paper so that it forms a cone (see the picture at the right.) and moisten it with a little distilled water so it will stick to the funnel.

To carry out the filtration, stir the solution to be filtered to suspend to solid substance. Then pour the mixture into the funnel as shown in the figure and allow it to filter into the beaker below. Use a little distilled water to transfer the solid to the filter paper. Do not discard the filtrate

7. After all the solid is on the filter paper and the liquid has drained into a beaker, carefully remove the filter paper and put it in the oven to dry. You will mass it the next day.

8. Isolate the sodium chloride from its solution by evaporating it in an evaporating dish and watch glass of known mass. You have already done a procedure like this so you should know how to do it without specific instructions. (A picture is provided to remind you). If you have a lot of solution, only fill the evaporating dish 1 2

2 3to full and add solution as it boils away.

9. Make certain that the solid product is dry when you mass it.

1st DAY CLEAN UP

10. Wash the stirring rod, large beakers, the evaporating dish and the watch glass with tap water and a scrubbing sponge. Rinse them, along with the funnel, thoroughly. Put the small beakers in the bus pan at the front of the room. Wipe your counter top with a damp sponge.

2nd DAY PROCEDURE and CLEAN UP 11. Mass the filter paper and precipitate. CaCO3 is chalk, so it can go in the regular wastebasket.

ANALYSIS AND CONCLUSIONS: 1.a. Using the number of moles of either reactant, determine the number of moles of CaCO3 that

should be produced in your experiment. b. Calculate the theoretical yield of CaCO3 in grams. c. Calculate the actual yield of CaCO3 in grams. d. Calculate the percent yield of CaCO3 in the experiment. e. Comment on one specific experimental error that could explain the disagreement between

the actual and theoretical yields of CaCO3. 2a-e. Answer questions 1a-e for NaCl. 3. The products of the reaction, CaCO3 and NaCl, are both white solids. Explain how we were able to separate them from each other despite their similarities in physical appearance.

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Lab 4-3: Titration of Vitamin C in Fruit Juice

(INDIVIDUAL) BACKGROUND Ascorbic acid, commonly known as vitamin C, is a water soluble vitamin, essential to human health. Vitamin C plays an important part in the formation of healthy teeth, red blood cells, and various connective tissues. Wounds also heal very slowly if vitamin C is absent from a person’s diet. Since vitamin C is water soluble, your body cannot store it for later use. It must be consumed every day. An average high school student needs 100 mg (0.100 g) of vitamin C a day to remain healthy. In this exercise you will have the opportunity to determine the volumes of your favorite beverages that are required to provide you with the minimum daily requirement of this vitamin. Vitamin C is not very stable. It reacts readily with oxygen and should therefore, not be left open to the air. Also, it can not withstand exposure to light or heat for extended periods of time. Ascorbic acid reacts with iodine in an acidic aqueous solution according to the following skeletal redox reaction: I2(aq) + C6H8O6(aq) → C6H6O6(aq) + I-

(aq)

Iodine + ascorbic acid → dehydroascorbic acid + Iodide ion (orange) Vitamin C (colorless) An aqueous iodine solution is a deep orange/red color, and solutions containing iodide ions are colorless. When an iodine solution is added to a solution containing an excess of ascorbic acid, the orange/red color disappears. EXPERIMENTAL METHOD

You will determine (1) the concentration of the iodine solution and (2) the mass of ascorbic acid in vitamin C-containing beverages by making use of the chemistry described above. The technique you will use to add the iodine solution to the ascorbic acid solution and the beverages is called a titration. This technique will be described in detail in the procedure section and demonstrated in class. You will add iodine solution to the ascorbic acid solution or beverage until the ascorbic acid has been completely consumed. How will you know that this particular “end point” has occurred? You will add some “starch indicator” to the solution of ascorbic acid or beverage. In this reaction, I2 reacts preferentially with C6H8O6 (Vitamin C). When all of the Vitamin C has reacted, however, any additional I2 reacts with the starch indicator, turning the solution dark blue. Thus, the exact moment the solution turns dark blue, you will know that all of the Vitamin C has been used up!

Buret filled

with I2

Solution containing Vitamin C

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PROCEDURE PART 1 - STANDARDIZATION OF THE IODINE SOLUTION

In Part 1 you will measure the volume of an iodine solution of unknown concentration that is required to completely react with an ascorbic acid solution of known concentration and volume. This will allow you to calculate the concentration of the iodine solution. This is known as “standardizing” the iodine solution.

1. Pour 40-50 mL of the iodine solution into a small beaker.

2. Be sure the valve on the buret is closed (horizontal position), and then use a funnel to pour the solution into the buret. Be careful not to overfill the buret.

3. Position the beaker under the tip of the buret. Open the valve (vertical position) and quickly close it so that the iodine solution flows through the tip of the buret.

4. Record the initial reading of iodine in the buret. For example, IF the buret is full, the initial reading would be 0.00 mL. Note that the buret does NOT have to be full. You will be using the buret in another experiment this year; the procedure for filling it and recording data is always the same.

5. Record the concentration of the standard ascorbic acid solution. 6. Use the graduated cylinder to measure exactly 25.0 mL of the ascorbic acid

solution. Record this volume. You can use the dropper to remove solution from the graduated cylinder if you overfill it.

7. Pour the solution into a clean 125 mL Erlenmeyer flask. 8. Rinse the graduated cylinder with distilled water and add the rinsings to the

flask. 9. Add about 10 drops of starch indicator. 10. Put a magnetic stir bar in the solution, place it on the stage of the stirrer and adjust the stirrer

to a moderate rate. Avoid splashing the solution on the sides of the flask. 11. Be as accurate as possible as you add the iodine solution to the ascorbic acid solution.

THIS REQUIRES PATIENCE. There are two techniques you can choose from:

You can try to set the valve on your buret to deliver the iodine solution in a steady drop-by-drop fashion until the characteristic blue color of the endpoint persists. You must monitor this situation closely so that you do not go past the end point.

The more traditional technique is to begin by adding relatively large volumes of the iodine solution say 0.50 mL at a time and then decrease the volume of each addition as you move toward the end point. You will probably see the blue color of the I2-starch complex when the solutions initially come in contact, but the color will disappear as the solution is stirred. The longer the blue color persists, the closer you are to the end point. As you approach the end point you should decrease the volume of iodine solution that you adding until your final additions are literally one or two drops. Continue until the characteristic blue color does not disappear.

12. When you reach the end point record the final reading of the iodine solution in the buret. 13. Discard the solution in the flask into the sink. Rinse the flask and stir bar with tap water and

then distilled water. 14. Repeat the previous procedure with another 25.0 mL portion of the standard ascorbic acid

solution. You should know the approximate end point volume of iodine solution from your first trial so the second trial should go quicker, but be patient and as accurate as possible as you near the end point.

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PROCEDURE PART 2 - TITRATION OF BEVERAGES CONTAINING VITAMIN C

In Part 2 you will measure the volume of the iodine solution that is required to completely react with the ascorbic acid in a known volume of beverage. Remember, the data from Part 1 will allow you to determine the concentration of the iodine solution. If you know the volume and concentration of the iodine solution required to completely react with the ascorbic acid in a known volume of beverage, you can calculate the moles and mass of ascorbic acid in that measured volume of beverage. This will allow you to determine the volume of that beverage that is needed to supply the minimum daily requirement of vitamin C. You may want to add more iodine solution to your buret. The procedure is identical to the procedure in Part 1 except that you use the vitamin C-containing beverages that you and your classmates brought to lab instead of the standard ascorbic acid solution. 15. Record the brand name of each beverage you test. 16. Measure out 25.0 mL of the vitamin C-containing beverage and pour it into the 125 mL

Erlenmeyer flask. If you use the dropper, rinse it between drinks. 17. Add 10 drops starch indicator. 18. Record the initial reading of iodine solution. 18. Titrate the beverage with iodine solution until the endpoint is reached. 19. Record the final reading of iodine solution in the buret. Repeat this procedure with one other beverage. Be sure to rinse everything between beverages. CLEAN UP - 20. When you have finished, drain any extra iodine solution back into the stock bottle. Put the lids on the stock bottles and tighten them. 21. Make certain that the buret has been thoroughly rinsed in the sink. Allow several mL of water to run through the buret tip. 22. Store the buret EXACTLY as you found it (upside down in the clamp - valve open). 23. Rinse all of your equipment thoroughly with tap water. 24. Rinse the dropper by drawing tap water inside and swirling it. 25. Invert the flask, beaker and grad in the rack to dry. Put the dropper and the funnel in the rack as well. 26. Place the magnetic stir bar on the stage of the stirrer. 27. Clean up any spills on the counter with a damp sponge. Your teacher will tell you what to do with any left over beverages.

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ANALYSIS AND CONCLUSIONS: 1. Balance the following “non-trivial” redox reaction in acidic aqueous solution. Show both balanced half reactions as well as your final answer.

I2 + C6H8O6 → C6H6O6 + I-1

PART 1: CALCULATING THE CONCENTRATION OF THE IODINE SOLUTION 2.a. Calculate the volume of the iodine solution added in each titration. b. Since you did the titration in Part 1 twice, average the two volumes of iodine calculated in

question 2a. 3.a. Calculate the number of moles of ascorbic acid in the original sample of the standard

ascorbic acid solution. b. Use stoichiometry to determine the number of moles of iodine that needed to be added to

reach the end point. c. Calculate the concentration (molarity) of the iodine solution.

PART 2 CALCULATING THE MASS OF VITAMIN C IN EACH BEVERAGE SAMPLE

4. Answer the following questions for one of the two beverages. a. Calculate the number of moles of iodine added from the buret. b. Use stoichiometry to determine the number of moles of vitamin C (ascorbic acid) in the

beverage sample. c. Calculate the mass (in grams) of vitamin C in the beverage sample that was titrated. d. Convert the mass of vitamin C to milligrams (mg). e. Recall that only 25.0 mL of the beverage was titrated in this experiment. Calculate the

volume of the beverage you would have to drink to fulfill your daily requirement for vitamin C. Look back at the information in the first paragraph of this handout and use a ratio.

5. Repeat calculations 4a-e for the second beverage. 6. Which beverage has a higher vitamin C content? 7. If you were to repeat Part 2 of the experiment using a more concentrated solution of iodine, will you need to add a larger or a smaller volume of iodine to reach the end point for each beverage? Explain briefly.

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Unit 4 Practice Test

Part 1: Multiple Choice

1. A 0.10 M CaCl2 solution is prepared by a. dissolving 0.10 moles of CaCl2 in 1 liter of water b. dissolving 0.10 moles of CaCl2 in enough water to make a 1 liter solution c. dissolving 0.10 moles of CaCl2 in 1 kilogram of water d. dissolving 0.10 grams of CaCl2 in one liter of water 2. What is the concentration of 10 mL of 0.25 M NaHCO3? a. 2.5 moles/liter b. 40 moles/liter c. 0.25 moles/liter d. 0.0025 moles/liter Use the following balanced equation to answer the next three questions (3-5) Zn(s) + 2 HCl(aq) → ZnCl2(aq) + H2(g)

3. How many moles of HCl are required to react with 0.311 moles of zinc? a. 2 b. 0.311 c. 0.622 d. 0.155 4. If a small amount of zinc remains at the conclusion of a reaction between zinc and hydrochloric acid, which reagent is in excess? a. zinc b. hydrochloric acid c. neither d. not enough information has been given 5. When a student combined 2.0 moles of HCl with an excess of zinc, 0.9 moles of hydrogen gas was produced. What was the percent yield of hydrogen gas? a. 90% b. 45% c. 18% d. 100%

6. What is the concentration of sulfate ions in a solution of 3.0 M Al2(SO4)3? a. 0.30 M b. 0.10 M c. 9.0 M d. 3.6 M

7. What is the molarity of a solution made by mixing 125 mL of 12.0 M HCl with 375 mL water?

a. 4.00 M b. 9.00 M c. 36.0 M d. 3.00 M

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Part 2: Problems

You must: SHOW ALL WORK for any answer that requires a calculation. Make certain that each numerical answer has a UNIT. Make certain that each numerical answer is expressed with the correct number of

SIGNIFICANT DIGITS.

8. A student evaporated the water in 0.200 liter of 0.350M AgNO3. How many grams of silver nitrate was he able to retrieve from this solution?

9. How would you prepare 500.0 mL of a 2.0 M CuSO4 solution? (i.e. what mass of solute and how much water?)

10. How would you prepare 750.0 mL of a 2.60 M H2SO4 solution from a 15.0 M solution of H2SO4? (i.e. what volume of concentrated solution and what volume of water?)

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(11-13) Use the following balanced equation and the accompanying list of molar masses to answer the next set of questions.

3 MgO(s) + 2 Fe(s) → Fe2O3(s) + 3 Mg(s) Substance MM (g/mole) MgO 40.3 11. How many moles of iron(III) oxide can be produced Fe 55.8 by the reaction of 5.0 moles of iron with an excess of Fe2O3 159.6 magnesium oxide? Mg 24.3 12. How many grams of iron are required to produce 7.29 g of magnesium by this reaction? 13. A chemist mixed 25.0 grams of MgO with 25.0 grams of Fe. Answer the following questions about this combination of reactants. a. Which is the limiting reactant? (Show work to support your answer.) b. How many moles of iron (III) oxide will be produced? c. How many grams of each reactant will be “left over” after the reaction is complete?

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14. A magnesium rod is placed in 850 mL of 0.500 M AgNO3. When the reaction is complete, the remaining portion of the rod is removed.

a. Write the balanced molecular equation for the reaction that occurs. _____________________________________________________________________________ b. What is the theoretical yield of silver in this reaction? c. When a student did this experiment, she actually obtained 41.2 grams of silver. What was her percent yield?

15.a. Predict the products and write a balanced molecular equation for the following neutralization reaction:

______ H3PO4 + ______ Ba(OH)2 _________________________________

b. 34.8 mL of 0.150 M H3PO4 is needed to completely neutralize 30.0 mL of a Ba(OH)2 solution of unknown concentration. Calculate the concentration of the Ba(OH)2 solution.

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Part 3: Essays

There may be one or more short essay questions on the test. The question(s) will be taken from the following topics unless you are notified otherwise.

E1. Explain how you would make a desired volume of an aqueous solution of accurate concentration using a solid solute and water or a concentrated solution and water.

E2. You did an experiment in which you carried out the following reaction: CaCl2(aq) + Na2CO3(aq) → CaCO3(s) + 2 NaCl(aq)

The purpose of the lab was to synthesize and isolate calcium carbonate and sodium chloride and calculate the percent yield of these products. Discuss how given technique mistakes would affect the results of this lab.

E3. You did an experiment in which you performed a titration of ascorbic acid with an iodine solution (using a starch indicator). Discuss the general experimental techniques and analytical techniques involved in this lab. Unit 3 Practice Test Answers 1. b 2. c 3. c 4. a 5. a 6. c 7. d 8. 11.9 g 9. 159 g CuSO4 and enough water to make 500 mL of solution 10. 0.130 L (or 130 mL) of 15.0 M H2SO4 and .620 L (or 620 mL) of water 11. 2.5 mol Fe2O3

12. 11.2 g Fe

13. a. MgO is limiting reactant b. 0.207 mol Fe2O3 c. MgO none leftover 2.0 g Fe leftover 14. a. 2 AgNO3 + Mg → Mg(NO3)2 + 2 Ag b. 46 g Ag c. 90% 15.a. 2 H3PO4 + 3 Ba(OH)2 Ba3(PO4)2 + 6 H2O b. 0.261 M

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Answers to Unit 4 Homework

Balanced equations and short answers

Assignment 1 p.3-5 problems 4-9

6.a. 3 Zn + 2 H3PO4 → Zn3(PO4)2 + 3 H2

7.a. 2 H2O → 2H2 + O2

8.a. 4 Fe + 3 O2 → 2 Fe2O3 Assignment 2 p.6 problem 11

11.a. C u + Cl2 → CuCl2 Assignment 3 p.10 problem 13 13.d. The yield of NO2 would increase because NO is the limiting reactant and additional NO

would react with the excess O2. e. The yield of NO2 would remain the same because O2 is already in excess and there isn't any

NO for the additional O2 to react with. Assignment 7 p.19-20 problems 24-28 26.a. 2 NaOH(aq) + H2CO3(aq) 2 H2O(l) + Na2CO3(aq)

27.a. 2 NO3- + 6 Cl- + 8 H+ → 2 NO + 3 Cl2 + 4 H2O

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