Materials Included In Kit

Additional Materials Required

Prelab Preparation

Hydrochloric acid solution, 4 M: Dilute 167 mL of 6 M hydrochloric acid to a final volume of 250 mL with distilled or deionized water. Always add acid to water, and mix well prior to dispensing.

Safety Precautions

Hydrochloric acid is corrosive to skin and eyes and toxic by inhalation or skin absorption. Avoid contact with eyes and skin and keep acid neutralizer on hand to clean up spills. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. For the gas collection experiment, do not use more than 0.5 g of calcium carbonate. The concentration of hydrochloric acid must not exceed 6 M in any experiment. Remind students to wash their hands thoroughly with soap and water before leaving the laboratory. Please follow all laboratory safety guidelines. Always review current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Please consult your current Flinn Scientific Catalog and Reference Manual for general guidelines and specific procedures, and review all federal, state, and local regulations that may apply, before proceeding. Leftover mixtures from each kinetics trial and excess hydrochloric acid solutions may be neutralized with base according to Flinn Suggested Disposal Method #24b. Excess hydrochloric acid may also be saved for future use. Solid calcium carbonate (marble chips) may be packaged for landfill disposal according to Flinn Suggested Disposal Method #26a.

Lab Hints

• The Introductory Activity may be done as an individual experiment or an interactive classroom demonstration to encourage participation and discussion.
• Encourage student discussion of the optimum number of concentrations and trials for reliable results. The reaction should be carried out at a minimum of three different HCl concentrations—two values will always give a straight line relationship between reaction rate and concentration. Additional trials may be done either at different concentrations or at the same concentration to average the results.
• It may be helpful to lubricate the plunger of the syringe with silicone grease or petroleum jelly to reduce friction. Apply a small dab of grease to the black rubber gasket only.
• Many students are initially confused about the relationship between the time they measure for the reaction to occur and the rate or speed of the reaction. Using car travel as an analogy usually clears up the confusion pretty easily. Although individual rates may not be accurate, the trends in reaction rate as a function of concentration and temperature are reproducible.
• Several variables will affect the rate of reaction and must be controlled in order to isolate the effect of concentration. The amount of hydrochloric acid consumed in the reaction is twice the number of moles of calcium carbonate, which is 0.005 moles in the Introductory Activity. If the volume of 6 M hydrochloric acid used is 10 mL, the initial number of moles of HCl present is 0.06 moles, and the amount of HCl consumed is (0.010/0.060) x 100, or 17% of the total. This percentage is greater than the 5–10% reaction generally advised for the method of initial rates and means that the reactant concentration is not a controlled variable. At lower concentrations of acid, the consumption ratio is even more pronounced, leading to problems with the design of the experiment. Students may consider this factor and choose to increase the volume of solvent (but not necessarily the amount of calcium carbonate).
• The inquiry-based approach outlined in this experiment can be easily extended to study other factors that affect the rate of reaction. Similar experiments can be designed to investigate how the surface area of calcium carbonate or the temperature of the acid influences the reaction rate. Do not use an acid concentration > 2 M to investigate the effect of temperature on the reaction rate. Do not exceed a temperature of 50 °C.

Further Extensions

Opportunities for Undergraduate Research
Investigate the effect of other variables, such as temperature, particle size or surface area, and the presence of a catalyst, on the rate of decomposition of calcium carbonate with hydrochloric acid.

Answers to Prelab Questions

1. Collision theory states that the rate of a reaction depends on the number of collisions between molecules or ions, the average energy of the collisions and their effectiveness. Does the general effect of concentration on reaction rate support the collision theory? Explain.

   In general, increasing the concentration of reactants also increases the rate of a reaction. This is consistent with collision theory because increasing the number of molecules or ions present in solution will increase the number of collisions between molecules. Concentration does not affect the energy or effectiveness of collisions.

2. The reaction of solid calcium carbonate with hydrochloric acid is a heterogeneous reaction. The rate law for this reaction will have the following form: rate = k[HCl]ⁿ. Explain why the concentration of calcium carbonate does not appear in the rate law.

   The concentration of a pure solid is equal to the density of the substance divided by its molar mass. It is an intensive property that is independent of the mass of the solid and also a constant that does not change much with temperature. The concentration of a solid does not change during the course of a reaction and thus does not appear in the rate law or equilibrium constant expressions for a reaction.

3. Read the entire Procedure for the Introductory Activity. In step 5, why is it necessary to make sure the stopcock is in the open position and then immediately replace the stopper and syringe assembly in the Erlenmeyer flask after adding the hydrochloric acid to the marble chips?

   Adding acid to the calcium carbonate will result in the immediate evolution of carbon dioxide gas, which will escape the container if the flask is not immediately stoppered. This will reduce the accuracy of the gas volume measurements that are used to calculate the reaction rate. The stopcock must be in the open position before the gas collection assembly is placed in the flask. Generating a gas in a closed system is dangerous and may cause the stopper or syringe to be blown out of the system, spraying caustic hydrochloric acid everywhere.

4. The average rate of reaction of hydrogen peroxide with iodide ions to produce iodine was determined for three initial concentrations of hydrogen peroxide as shown in the following table. What is the order of the reaction with respect to hydrogen peroxide? Explain your reasoning.
   {12751_PreLabAnswers_Table_2}
   Doubling the reactant concentration from 0.022 M to 0.044 M doubled the reaction rate. This is consistent with a first order reaction in the rate law. The conclusion is corroborated by comparing a second pair of results, concentrations of 0.044 M and 0.089 M. Again, doubling the concentration caused the rate to increase by a factor of two.

Sample Data

Introductory Activity

Gas Collection and Mass Loss Methods Analysis and Conclusions
The rate of decomposition of calcium carbonate with hydrochloric acid depends on the concentration of acid. Two different methods for determining the rate of reaction were compared. In the gas collection method, 0.50 g of ground or powdered calcium carbonate was combined with 10.0 mL of hydrochloric acid. The quantities were doubled to 1.00 g calcium carbonate and 20.0 mL of acid when using the mass loss method. (The actual mass of HCl added to the flask is critical.) Three concentrations of acid were studied—1 M, 2 M and 4 M.

Rates of reaction were calculated using the method of initial rates by determining the slope for the linear portion of each curve shown in the graphs. Depending on the concentration of acid, the linear portion of the curve and the initial rate corresponded to approximately 10−30% of reaction completion after 3 minutes. The methods gave comparable results indicating that the rate is roughly proportional to the concentration of HCl, consistent with a first order reaction. More experiments are needed to verify this conclusion due to the experimental errors involved. Comparing the results obtained using two methods highlights the role and importance of experimental error. The gas collection method appeared to give more reliable results at lower concentrations of acid (1 M or 2 M), while the mass loss method appeared to give more consistent results with higher concentrations (2 M or 4 M). The rapid mass loss observed within one minute using 4 M HCl suggests that a significant amount of gas may have been lost in the gas collection experiment for this concentration of acid. The flask is open to the air for 10−20 seconds as the acid is added and the syringe assembly is attached. Conversely, the very low rate of gas collection when using 1 M HCl indicates that the mass loss method suffers due to a lack of precision in the measurements with a centigram (0.01-g) balance. The number of moles of HCl present in 10 or 20 mL of 1 M acid corresponds to the exact stoichiometric amount needed to react completely with 0.5 or 1 g of CaCO₃, respectively, depending on the method used. This is a another source of error, since the concentration decreases substantially as the reaction starts, further slowing down an already very slow reaction.

Answers to Questions

Guided-Inquiry Design and Procedure

1. Analyze the graph from the Introductory Activity. Does the amount of CO₂ increase linearly with time, or does it level off as the reaction proceeds? Explain the shape of the curve based on the rate of the reaction and the concentration of HCl versus time.

   The amount of CO₂ generated increases very rapidly and in a linear fashion at first but then begins to level off or hit a plateau after about four minutes (depending on the trial). The slope of the linear portion of the graph is equal to the reaction rate. Since rate is proportional to the concentration of reactant(s), and the concentration of the reactant(s) decreases as the reaction proceeds, the rate levels off as more and more of the reactant(s) are consumed.

2. Initial rates are generally used to compare reaction rates for different concentrations of reactants. The initial rate is calculated from the slope or linear portion of the graph of the amount of product versus time. Estimate the initial rate for the reaction of CaCO₃ with HCl, and express in units of volume of CO₂ per minute.

   The initial rate of reaction was almost 25 mL of CO₂ per minute. This is the average of two trials. The reaction was very rapid with 6 M HCl, leading to large potential losses of CO₂ in the interval between adding the acid and securing the gas collection assembly in the flask.

3. Calculate the number of moles of CaCO₃ and HCl used in the Introductory Activity. Determine the limiting reactant and use the ideal gas law to estimate the maximum volume of CO₂ that could be produced. Use the average room temperature and barometric pressure in the calculations.

   Molar mass of CaCO₃ = 100 g/mole
   Moles of CaCO₃ used = 0.50 g/100 g/mole = 5 x 10^–3 moles
   Moles of HCl = 6 moles/L x 0.010 L = 6 x 10^–2 moles
   Given the mole ratio of 2 moles HCl per mole CaCO₃, the reaction mixture contains a six-fold molar excess of HCl. The limiting reagent is CaCO₃ and the theoretical yield of CO₂ is 5 x 10^–3 moles. Substituting this value of n in the ideal gas law (PV = nRT), and assuming the temperature was 22 °C (295 K) and the pressure 740 mmHg (0.974 atm), the theoretical or maximum volume of CO₂ that could be collected is 124 mL.

4. Is the volume of the syringe sufficient to contain all of the CO₂ that could be produced? What was the average percent of reaction completion after 10 minutes? Explain in terms of potential sources of error in the experiment.

   The volume of the syringe is greater than the theoretical yield of CO₂ gas that could be produced. The average percent of reaction completion was about 90% after 10 minutes, suggesting that some carbon dioxide was unaccounted for, either due to the solubility of carbon dioxide gas in the water (reaction mixture) or gas escaping to the atmosphere in the time between adding the acid and securing the gas collection apparatus in the flask. Since the reaction with 6 M HCl is very rapid, the effect of the latter error will be less pronounced when lower concentrations of HCl are used.

5. Two alternative procedures may be used to compare the effect of concentration on reaction rate. The first was demonstrated in the Introductory Activity. The second procedure involves the change in mass of the reaction mixture versus time. How will the mass of the system change as the reaction proceeds? What quantity will be proportional to the amount of CO₂ produced?

   The mass of the reaction mixture (assumed to be the total mass of CaCO₃ and added HCl) will decrease as the reaction proceeds due to the loss of CO₂ to the atmosphere. The difference between the initial combined mass and the mass at any time should equal the mass of CO₂ produced, since all other products remain in the flask or beaker. Thus, graphing mass loss versus time should have the same general shape or type of curve as seen in the Introductory Activity, and the initial rate of the reaction can be obtained from the slope of the linear portion of the curve.

6. What is the theoretical mass of CO₂ that can be produced from (a) 0.50 g of CaCO₃ and (b) 1.00 g of CaCO₃? Which reactant quantity is more suitable for the alternative procedure? Explain based on the precision of the balance and other factors and discuss how this choice would affect the volume of HCl that is used.

   Molar mass of CO₂ = 44 g/mole
   Theoretical number of moles of product = 5 x 10^–3 moles if 0.50 g of CaCO₃, are used, and 1 x 10^–2 moles if 1.00 g of CaCO₃ are used. Predicted mass loss when the reaction is 100% complete:
   (a) 5 x 10^–3 moles x 44 g/mole = 0.22 g
   (b) 1 x 10^–2 moles x 44 g/mole = 0.44 g
   Option b (1.00 g CaCO₃ and 20 mL of HCl) is a better option for using a 0.01-g precision balance to determine the rate of reaction by measuring mass loss. Otherwise the mass loss will be very small, especially in the beginning of the reaction, which is the time period used in calculating the initial rate. A 0.001-g balance is preferred.

7. Identify the measurements that must be made for both procedures to investigate the effect of HCl concentration on the reaction rate and to determine the reaction order with respect to HCl. Name the independent and dependent variables for each series of experiments, and choose some suitable values for the independent variable.

   Reaction rates are calculated using the slope of the linear portion of the curve of amount of CO₂ versus time. The amount of CO₂ may be measured by volume (gas collection procedure) or mass (mass loss procedure). The concentration of HCl is the independent variable in experiments to determine the reaction order with respect to HCl, and the amount of gas produced is the dependent variable. Suitable concentrations of HCl for the independent variable are 1, 2, 3 and 4 M. The 6 M reaction is very fast and does not give reproducible results. The 1 M reaction may be very slow, but only needs to be followed to 10–20 % completion. If using the mass loss procedure, it is very important that the actual mass of HCl added to the flask be accurate. We recommend measuring the mass of a graduated cylinder plus acid, and then subtracting the mass of the cylinder again after adding the acid.

8. Discuss how the size or surface area of the marble chips might affect the purpose or design of the experiments. What is the best way to control this variable so that it remains constant and does not affect the analysis?

   Marble chips range in size and also mass so it is hard to get consistent amounts in different trials. The surface area of a solid will have a profound effect on the rate of a heterogeneous reaction and this variable must be controlled in order to investigate the effect of HCl concentration. This may be accomplished by grinding the marble chips with a mortar and pestle to a powder, thoroughly mixing the powder in a vial, and then removing different samples for the various trials.

Post-Laboratory Review

1. An unknown Group 1 metal carbonate M₂CO₃ (M = Li, Na or K) was reacted with excess 2 M HCl and the mass of CO₂ released was determined by mass difference. The initial mass of solid M₂CO₃ was 2. 002 g and the mass of CO₂ released was 1.206 g.
   1. Write the balanced chemical equation for the reaction of M₂CO₃ with HCl.

      M₂CO₃(s) + 2HCl(aq) → 2MCl(aq) + CO₂(g) + H₂O(l)

   2. What is the mole ratio of CO₂ to M₂CO₃?

      One mole of CO₂ is produced per mole of M₂CO₃.

   3. Calculate the molar mass of the unknown metal carbonate and identify the Group 1 metal.

      The moles of CO₂ produced is 1.206 g/44.0 g/mole = 0.0274 moles.
      The molar mass of M₂CO₃ is 2.002 g/0.0274 moles = 73.0 g/mole.
      Theoretical molar masses of M₂CO₃ are Li₂CO₃, 73.88 g/mole; Na₂CO₃, 106 g/mole; and K₂CO₃, 138.2 g/mole.
      The unknown metal is lithium.

2. The rate of reaction of 0.030 g of magnesium ribbon with 1 M hydrochloric acid was studied at four different temperatures by measuring the time required for the magnesium metal to disappear. The following data was recorded:
   {12751_Answers_Table_4}
   1. Calculate the number of moles of magnesium that reacted and the average reaction rate for each temperature.

      Moles of Mg reacted = 0.030 g/24.3 g/mole = 1.2 x 10⁻³ moles
      Sample calculation for average reaction rate at 2 °C: 1.2 x 10⁻³ moles/204 sec = 5.9 x 10⁻⁶ moles/sec

   2. Convert each temperature to kelvins and plot the average reaction time versus temperature in the following graph. Predict how long the reaction would take at 75 °C.
      {12751_Answers_Figure_4}
      At 75 °C = 348 K, the predicted reaction time would be very fast, perhaps about 10 sec.
   3. Using kinetic molecular theory and collision theory, explain why the absolute temperature scale (kelvins) is more appropriate than Celsius for explaining the effect of temperature on reaction rate. Does the effect of temperature on reaction rate support the collision theory of chemical reactions?

      Kelvin is more appropriate than Celsius for explaining the effect of temperature on reaction rate because the kinetic energy of molecules in motion is proportional to the absolute temperature in Kelvin. According to collision theory, the reaction rate depends on the number of collisions between molecules. As the absolute temperature increases, the kinetic energy of colliding molecules increases, leading to both a greater number collisions and a greater number of molecules with sufficient energy to overcome the activation energy barrier for the reaction.

Introduction

What factors determine how fast a chemical reaction will occur? The answer has applications in chemistry, food science, geology, ecology, and even art and architecture. Consider the weathering of beautiful marble statues from antiquity. The history of our civilization is gradually being eroded as acid in the environment dissolves the calcium carbonate in marble. Investigate the rate of decomposition of calcium carbonate with different concentrations of hydrochloric acid to learn more about kinetics and the rates of chemical reactions.

Concepts

• Kinetics
• Rate of reaction
• Collision theory
• Rate law
• Order of reaction
• Gas laws
• Acids and bases
• Neutralization
• Environmental chemistry

Background

Calcium carbonate, CaCO₃, is one of the most abundant minerals on the Earth. More than 4% of the Earth’s crust is composed of calcium carbonate. It is a major component in limestone, marble, seashells, bedrock, etc. Limestone and marble have been among the most widely used building materials for more than 5000 years, from the pyramids in Egypt to the Parthenon in Greece and the Taj Mahal in India. In many places, limestone is also the foundation of our Earth—literally, since it forms both bedrock and mountain ranges. Calcium carbonate dissolves in water to only a limited extent, but its solubility is greatly enhanced when the water is acidic. The gradual dissolution of marble and limestone, as well as coral and seashells, in acids is due to acid−base neutralization. The products of the neutralization reaction between calcium carbonate and hydrochloric acid are calcium chloride and carbonic acid, or H₂CO₃. Carbonic acid is unstable, decomposing to give carbon dioxide gas and water.

The rate of the overall reaction (Equation 3) and its dependence on the concentration of HCl are important concerns in environmental chemistry due to the combined effects of acid rain and ocean acidification. Kinetics is the study of the rates of chemical reactions. As reactants are transformed into products in a chemical reaction, the amount of reactants will decrease and the amount of products will increase. The rate of the reaction can be determined by measuring the amounts or concentrations of reactants or products as a function of time. In some cases, it is possible to use a simple visual clue to determine a reaction rate. Some of the “clues” that may be followed to measure a reaction rate include color intensity, amount of precipitate that forms, or amount of gas generated. In the case of the reaction of CaCO₃ with HCl, one of the products is a gas. Since either volume or mass of the gas is proportional to moles, the rate can be followed by measuring the time it takes for a specific volume or mass of carbon dioxide to be released. The reaction rate is calculated by dividing the quantity of carbon dioxide produced by the time. The rate of a reaction describes how fast the reaction occurs—the faster the rate, the less time that is needed for a specific amount of reactants to be converted to products. Some factors that affect the rates of chemical reactions include the nature of the reactants, their concentration, the reaction temperature, the surface area of solids, and the presence of catalysts. The relationship between the rate of a reaction and the concentration of reactants is expressed in a mathematical equation called a rate law. For a general reaction of the form:

A + B → C

the rate law can be written as:

Rate = k[A]ⁿ[B]^m

where k is the rate constant, [A] and [B] are the molar concentrations of the reactants, and n and m are exponents that define how the rate depends on the individual reactant concentrations. The rate decreases over the course of the reaction as the concentrations of reactants decrease. Rate laws are usually determined by analyzing the rate after approximately 10−20% of reactant(s) have been consumed.

The exponents n and m are also referred to as the order of reaction with respect to each reactant. In this example, the reaction is said to be nth order in A and mth order in B. In general, n and m will be positive whole numbers—typical values of n and m are 0, 1 and 2. When n = 0, the rate does not depend on the concentration of the reactant. When n = 1, the reaction will occur twice as fast when the reactant concentration is doubled, and when n = 2, the rate will increase by a factor of four when the reactant concentration is doubled. The values of the exponents must be determined by experiment—they cannot be predicted simply by looking at the balanced chemical equation.

Experiment Overview

The purpose of this inquiry lab is to design kinetics experiments for the heterogeneous reaction of calcium carbonate with hydrochloric acid. The investigation begins with an introductory activity to observe the evolution of carbon dioxide gas from the decomposition of calcium carbonate with acid. Special equipment is provided to collect and measure the volume of gas generated. The procedure provides a model for guided-inquiry design of experiments to determine the rate of reaction with different concentrations of acid. Using a cooperative approach, different groups will compare data for mass loss and volume of gas generation versus time. Initial rates and the rate law for the reaction are determined by graphical analysis of the results. The effects of temperature and particle size or surface area on the reaction rate provide additional opportunities for inquiry.

Materials

Prelab Questions

1. Collision theory states that the rate of a reaction depends on the number of collisions between molecules or ions, the average energy of the collisions and their effectiveness. Does the general effect of concentration on reaction rate support the collision theory? Explain.
2. The reaction of solid calcium carbonate with hydrochloric acid is a heterogeneous reaction. The rate law for this reaction will have the following form: rate = k[HCl]ⁿ. Explain why the concentration of calcium carbonate does not appear in the rate law.
3. Read the entire Procedure for the Introductory Activity. In step 5, why is it necessary to make sure the stopcock is in the open position and then immediately replace the stopper and syringe assembly in the Erlenmeyer flask after adding the hydrochloric acid to the marble chips?
4. The average rate of reaction of hydrogen peroxide with iodide ions to produce iodine was determined for three initial concentrations of hydrogen peroxide as shown in the following table. What is the order of the reaction with respect to hydrogen peroxide? Explain your reasoning.
   {12751_PreLab_Table_1}

Safety Precautions

Hydrochloric acid is corrosive to skin and eyes and toxic by inhalation or skin absorption. Avoid contact with eyes and skin and clean up all spills immediately. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron. For the gas collection experiment, do not use more than 0.5 g of calcium carbonate. The concentration of hydrochloric acid must not exceed 6 M in any experiment. Wash hands thoroughly with soap and water before leaving the laboratory. Please follow all laboratory safety guidelines.

Procedure

Introductory Activity

1. Read the entire Procedure before beginning.
2. Obtain about 0.5 g of calcium carbonate (marble chips) and measure the precise mass. Place the marble chips into a 125-mL Erlenmeyer flask.
3. Set up the gas-collection apparatus as shown in Figure 1. Make sure the rubber stopper fits securely in the flask and the stopcock is open. {12751_Procedure_Figure_1}
4. Carefully measure 10 mL of 6 M hydrochloric acid in a 10-mL graduated cylinder.
5. Remove the stopper and syringe assembly from the Erlenmeyer flask and quickly but carefully add the acid to the flask. Immediately replace the stopper and syringe assembly in the flask and start timing.
6. The syringe plunger will gradually expand out or lift up as gas is generated.
7. Measure the volume of gas in the syringe at one-minute intervals for 10 minutes. To overcome friction or resistance in the syringe, gently depress the plunger and then release it just before measuring the volume every minute.
8. Graph the volume (mL) of gas produced on the y-axis versus time (minutes) on the x-axis.

Guided-Inquiry Design and Procedure
Form a working group with other students and discuss the following questions.

1. Analyze the graph from the Introductory Activity. Does the amount of CO₂ increase linearly with time, or does it level off as the reaction proceeds? Explain the shape of the curve based on the rate of the reaction and concentration of HCl versus time.
2. Initial rates are generally used to compare reaction rates for different concentrations of reactants. The initial rate is calculated from the slope or linear portion of the graph of the amount of product versus time. Estimate the initial rate for the reaction of CaCO₃ with HCl, and express in units of volume of CO₂ per minute.
3. Calculate the number of moles of CaCO₃ and HCl used in the Introductory Activity. Determine the limiting reactant and use the ideal gas law to estimate the maximum volume of CO₂ that could be produced. Use the average room temperature and barometric pressure in the calculations.
4. Is the volume of the syringe sufficient to contain all of the CO₂ that could be produced? What was the average percent of reaction completion after 10 minutes? Explain in terms of potential sources of error in the experiment.
5. Two alternative procedures may be used to compare the effect of concentration on reaction rate. The first was demonstrated in the Introductory Activity. The second procedure involves the change in mass of the reaction mixture versus time. How will the mass of the system change as the reaction proceeds? What quantity will be proportional to the amount of CO₂ produced?
6. What is the theoretical mass of CO₂ that can be produced from (a) 0.50 g of CaCO₃ and (b) 1.00 g of CaCO₃? Which reactant quantity is more suitable for the mass loss procedure? Explain based on the precision of the balance and other factors, and discuss how this choice would affect the volume of HCl that is used.
7. Identify the measurements that must be made for both procedures to investigate the effect of HCl concentration on the reaction rate and to determine the reaction order with respect to HCl. Name the independent and dependent variables for each series of experiments, and choose some suitable values for the independent variable.
8. Discuss how the size or surface area of the marble chips might affect the purpose or design of the experiments. What is the best way to control this variable so that it remains constant and does not affect the analysis?
9. Write a step-by-step procedure for two alternative series of experiments to investigate the effect of HCl concentration on the reaction of CaCO₃ with HCl. Include the materials, glassware and equipment that will be needed, required safety precautions, concentrations and amounts of reactants, etc.
10. Review additional variables that may affect the accuracy or reproducibility of the experiments.
11. Carry out the experiment and record results in an appropriate data table.

Analyze the Results
Graph the results for each trial and calculate the average or initial rate for each concentration of HCl. Plot or analyze the rate versus HCl concentration to determine the reaction order. Compare and contrast the results obtained using the two alternative procedures and discuss any discrepancies as well as potential sources of experimental error in each procedure.

Student Worksheet PDF

12751_Student1.pdf