Reactions proceed faster at higher temperatures because increasing the temperature also increases the average kinetic energy of the reactant molecules, making it more likely that a collision between two reactant molecules will have sufficient energy to initiate the reaction. An old rule of thumb states that increasing the temperature by 10°C doubles the rate of reaction. (So, of course, reducing the temperature by 10°C halves the reaction rate.)
We’ll test that rule of thumb in this lab by reacting Alka-Seltzer tablets with water at different temperatures. When an Alka-Seltzer tablet is dropped into water, it emits carbon dioxide as it dissolves and its components react. We could get basic data about the reaction rate simply by determining how long it takes a tablet to dissolve completely in water at various temperatures. But using that method gives us data only about the starting and ending points. It doesn’t tell us, other than in the most general terms, whether the reaction rate is linear or if it changes over time.
We know that this reaction evolves gaseous carbon dioxide, which means that the mass of the reaction vessel and contents decreases as the reaction proceeds. By recording the mass of the reaction vessel and its contents periodically as the reaction progresses, we can gather the data needed to determine whether the reaction rate is linear and whether reaction temperature affects the linearity of the reaction rate.
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Weigh one of the tablets to 0.01 g and record its mass on line A of Table 12-1.
Add about 100 mL of cold tap water to the foam cup. (Use a foam cup rather than a beaker to keep the mass of the reaction vessel and its contents under the maximum capacity of your balance. If the combined mass of the foam cup, Alka-Seltzer tablet, and 100 mL of water is greater than the maximum capacity of your balance, reduce the quantity of water accordingly.)
Weigh the foam cup and water and record the mass on line B of Table 12-1.
Measure and record the temperature of the water on line C of Table 12-1.
With the cup and water still on the balance, drop the tablet into the cup.
Note the combined mass of the cup, water, and tablet every five seconds and record each mass in Table 12-1. (It may be helpful to have one person watching the clock while another calls out the mass reading at each 5-second milestone.)
Continue recording the changing mass until you reach one minute or until the reaction completes, as evidenced by the cessation of bubbling.
When the reaction completes, record the final mass of the cup, water, and tablet on line Q of Table 12-1.
Dispose of the spent solution and rinse out the cup.
Repeat steps 1 through 9, using hot tap water.
Repeat steps 1 through 9, using a mixture of half hot and half cold tap water.
Although the chemicals used in this laboratory are not hazardous, it is good practice to wear splash goggles, gloves, and protective clothing at all times.
All of the solutions from this laboratory can be flushed down the drain with plenty of water.
Item | Trial A | Trial B | Trial C |
A. Mass of tablet | ________.___ g | ________.___ g | ________.___ g |
B. Mass of cup + water | ________.___ g | ________.___ g | ________.___ g |
C. Temperature of water | ________.__ °C | ________.__ °C | ________.__ °C |
D. Mass at 0:00 (A + B) | ________.___ g | ________.___ g | ________.___ g |
E. Mass at 0:05 | ________.___ g | ________.___ g | ________.___ g |
F. Mass at 0:10 | ________.___ g | ________.___ g | ________.___ g |
G. Mass at 0:15 | ________.___ g | ________.___ g | ________.___ g |
H. Mass at 0:20 | ________.___ g | ________.___ g | ________.___ g |
I. Mass at 0:25 | ________.___ g | ________.___ g | ________.___ g |
J. Mass at 0:30 | ________.___ g | ________.___ g | ________.___ g |
K. Mass at 0:35 | ________.___ g | ________.___ g | ________.___ g |
L. Mass at 0:40 | ________.___ g | ________.___ g | ________.___ g |
M. Mass at 0:45 | ________.___ g | ________.___ g | ________.___ g |
N. Mass at 0:50 | ________.___ g | ________.___ g | ________.___ g |
O. Mass at 0:55 | ________.___ g | ________.___ g | ________.___ g |
P. Mass at 1:00 | ________.___ g | ________.___ g | ________.___ g |
Q. Mass at completion of reaction | ________.___ g | ________.___ g | ________.___ g |
R. Mass loss (D – Q) | ________.___ g | ________.___ g | ________.___ g |
S. Mass loss percentage [(R/A) · 100] | ________.___ % | ________.___ % | ________.___ % |
Q: | Q1: What effect did you observe temperature to have on reaction rate? ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ |
Q: | Q2: Based on the data you recorded in Table 12-1, does the 10°C rule of thumb provide a reasonably close approximation of the observed reaction rates in this laboratory? How far do your data depart from the expected values based on the rule of thumb? ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ |
Q: | Q3: Based on the data you recorded in Table 12-1, does reaction rate appear to be approximately linear over time? If you noticed an increase or decrease in reaction rate over time, propose at least one possible explanation. ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ |
If at least one of the reactants is a solid, the reaction proceeds faster if the solid is finely divided, because the surface area is larger in a finely divided solid, exposing more of that reactant to the other reactant or reactants. For example, a 50-pound bag of flour is essentially inert, because the flour, although finely ground, exposes little of its surface area to the air. But that same amount of flour dispersed as airborne dust, if ignited by a spark, explodes with force sufficient to flatten a large building. (Military fuel-air explosives, or FAEs, use this principle by vaporizing a liquid fuel and detonating it. During Operation Desert Storm, British troops reportedly sent a flash-priority report of a nuclear detonation, mistaking the detonation of an FAE for a tactical nuke.)
In the last laboratory, we examined reaction rates of Alka-Seltzer tablets, keeping the initial surface area constant and varying the temperature. In this laboratory, we’ll explore the effect on reaction rates of varying the surface area at constant temperature.
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Weigh one of the tablets to 0.01 g and record its mass on line A of Table 12-2.
Add about 100 mL of cold tap water to the foam cup. Again, make sure the combined mass of the foam cup, Alka-Seltzer tablet, and 100 mL of water is less than the maximum capacity of your balance. If not, reduce the quantity of water accordingly.
Weigh the foam cup + water and record the mass on line B of Table 12-2.
Measure and record the temperature of the water on line C of Table 12-2.
With the cup and water still on the balance, drop the tablet into the cup.
Note the combined mass of the cup, water, and tablet every five seconds and record each mass in Table 12-2. (It may be helpful to have one person watching the clock while another calls out the mass reading at each 5-second milestone.)
Continue recording the changing mass until you reach one minute or until the reaction completes, as evidenced by the cessation of bubbling.
When the reaction completes, record the final mass of the cup, water, and tablet on line Q of Table 12-2.
Dispose of the spent solution and rinse out the cup.
Repeat steps 1 through 9, using a tablet that you have split into quarters.
Repeat steps 1 through 9, using a tablet that you have crushed into powder. (This reaction may be too fast to time accurately; just do your best.)
Although the chemicals used in this laboratory are not hazardous, it is good practice to wear splash goggles, gloves, and protective clothing at all times.
All of the solutions from this laboratory can be flushed down the drain with plenty of water.
Item | Trial A (solid) | Trial B (chunks) | Trial C (powder) |
A. Mass of tablet | ________.___ g | ________.___ g | ________.___ g |
B. Mass of cup + water | ________.___ g | ________.___ g | ________.___ g |
C. Temperature of water | ________.__ °C | ________.__ °C | ________.__ °C |
D. Mass at 0:00 (A + B) | ________.___ g | ________.___ g | ________.___ g |
E. Mass at 0:05 | ________.___ g | ________.___ g | ________.___ g |
F. Mass at 0:10 | ________.___ g | ________.___ g | ________.___ g |
G. Mass at 0:15 | ________.___ g | ________.___ g | ________.___ g |
H. Mass at 0:20 | ________.___ g | ________.___ g | ________.___ g |
I. Mass at 0:25 | ________.___ g | ________.___ g | ________.___ g |
J. Mass at 0:30 | ________.___ g | ________.___ g | ________.___ g |
K. Mass at 0:35 | ________.___ g | ________.___ g | ________.___ g |
L. Mass at 0:40 | ________.___ g | ________.___ g | ________.___ g |
M. Mass at 0:45 | ________.___ g | ________.___ g | ________.___ g |
N. Mass at 0:50 | ________.___ g | ________.___ g | ________.___ g |
O. Mass at 0:55 | ________.___ g | ________.___ g | ________.___ g |
P. Mass at 1:00 | ________.___ g | ________.___ g | ________.___ g |
Q. Mass at completion of reaction | ________.___ g | ________.___ g | ________.___ g |
R. Mass loss (D – Q) | ________.___ g | ________.___ g | ________.___ g |
S. Mass loss percentage [(R/A)·100] | ________.___ % | ________.___ % | ________.___ % |
Reactions proceed faster at higher concentrations because more reactant molecules are available and therefore collisions between reactant molecules are more likely. In the two preceding labs, we used Alka-Seltzer tablets and water to demonstrate the effects of temperature and surface area on reaction rates. Obviously, we’ll have to use some other method to demonstrate the effect of concentration on reaction rates, because the amounts of the citric acid and sodium bicarbonate reactants in a fizzy tablet are fixed and cannot be changed.
In this lab, we’ll use another OTC medicine with varying concentrations of hydrochloric acid to demonstrate the effect of concentration on reaction rates. Some (but not all) antacid tablets contain primarily calcium carbonate, along with flavoring, binders, and other inactive ingredients. These tablets neutralize excess stomach acid, which is actually dilute hydrochloric acid, according to the following equation:
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + CO2(g) + H2O(l)
We’ll use exactly this reaction to observe and quantify the effect of concentration of reaction rates.
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Label three foam cups 4 M, 2 M, and 1 M.
In the first foam cup, make up 100 mL of 4 M hydrochloric acid by adding 33.3 mL of 12 M HCl to 66.7 mL of water. Again, make sure that the combined mass of the foam cup, antacid tablet, and 100 mL of hydrochloric acid is less than the maximum capacity of your balance. If not, reduce the quantity of hydrochloric acid accordingly.)
In the second foam cup, make up 100 mL of 2 M hydrochloric acid by adding 16.7 mL of 12 M HCl to 83.3 mL of water.
In the third foam cup, make up 100 mL of 1 M hydrochloric acid by adding 8.3 mL of 12 M HCl to 91.7 mL of water.
Always add acid to water. Adding water to a concentrated acid can cause splattering. Diluting acid produces heat. Because we want concentration to be the only variable, it’s important to ensure that all three acid solutions are at the same initial temperature. Make your dilute acid solutions ahead of time and allow all three cups to cool to room temperature before proceeding.
Weigh one of the antacid tablets to 0.01 g and record its mass on line A of Table 12-3.
Place the first foam cup with 4 M hydrochloric acid on the balance and record the mass to 0.01 g on line B of Table 12-3.
Measure and record the temperature of the HCl solution on line C of Table 12-3.
With the cup and its contents still on the balance, drop the antacid tablet into the cup.
Note the combined mass of the cup, acid, and tablet every five seconds and record each mass in Table 12-3. (It may be helpful to have one person watching the clock while another calls out the mass reading at each 5-second milestone.)
Continue recording the changing mass until you reach one minute or until the reaction completes, as evidenced by the cessation of bubbling.
When the reaction completes, record the final mass of the cup, water, and tablet on line Q of Table 12-3.
Dispose of the spent solution and rinse out the cup.
Repeat steps 6 through 9, using 2 M hydrochloric acid.
Repeat steps 6 through 9, using 1 M hydrochloric acid.
Neutralize the spent acid solutions with sodium bicarbonate or another base and flush the neutralized solutions down the drain with plenty of water.
Item | Trial A (4 M HCl) | Trial B (2 M HCl) | Trial C (1 M HCl) |
A. Mass of tablet | ________.___ g | ________.___ g | ________.___ g |
B. Mass of cup + hydrochloric acid | ________.___ g | ________.___ g | ________.___ g |
C. Temperature of hydrochloric acid | ________.__ °C | ________.__ °C | ________.__ °C |
D. Mass at 0:00 (A + B) | ________.___ g | ________.___ g | ________.___ g |
E. Mass at 0:05 | ________.___ g | ________.___ g | ________.___ g |
F. Mass at 0:10 | ________.___ g | ________.___ g | ________.___ g |
G. Mass at 0:15 | ________.___ g | ________.___ g | ________.___ g |
H. Mass at 0:20 | ________.___ g | ________.___ g | ________.___ g |
I. Mass at 0:25 | ________.___ g | ________.___ g | ________.___ g |
J. Mass at 0:30 | ________.___ g | ________.___ g | ________.___ g |
K. Mass at 0:35 | ________.___ g | ________.___ g | ________.___ g |
L. Mass at 0:40 | ________.___ g | ________.___ g | ________.___ g |
M. Mass at 0:45 | ________.___ g | ________.___ g | ________.___ g |
N. Mass at 0:50 | ________.___ g | ________.___ g | ________.___ g |
O. Mass at 0:55 | ________.___ g | ________.___ g | ________.___ g |
P. Mass at 1:00 | ________.___ g | ________.___ g | ________.___ g |
Q. Mass at completion of reaction | ________.___ g | ________.___ g | ________.___ g |
R. Mass loss (D – Q) | ________.___ g | ________.___ g | ________.___ g |
S. Mass loss percentage [(R/A) · 100] | ________.___ % | ________.___ % | ________.___ % |
Q: | Q1: What effect did you observe concentration to have on reaction rate? ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ |
Q: | Q2: Based on the data you recorded in Table 12-3, is the effect of concentration on reaction rate linear? ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ |
Q: | Q3: Based on the data you recorded in Table 12-3, at any particular concentration does reaction rate appear to be approximately linear over time? If you noticed an increase or decrease in reaction rate over time, propose at least one possible explanation. ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ |
A catalyst is a substance that increases the rate of a chemical reaction, but is not consumed or changed by the reaction. A catalyst works by reducing the activation energy needed to initiate and sustain the reaction. For example, two molecules of hydrogen peroxide can react to form two molecules of water and one molecule of molecular oxygen gas by the following reaction:
2 H2O2(aq) → 2 H2O(l) + O2(g)
At room temperature, this reaction occurs very slowly because few of the collisions between hydrogen peroxide molecules have sufficient energy to activate the reaction. Furthermore, commercial hydrogen peroxide solutions, such as the 3% hydrogen peroxide solution sold in drugstores and the 6% solution sold by beautician supply stores, are treated with stabilizers (sometimes called negative catalysts) that increase the activation energy for the reaction, further inhibiting it from occurring.
If you add a catalyst to a solution of hydrogen peroxide, the effect is immediately evident. The solution begins bubbling, as oxygen gas is evolved. Numerous substances can catalyze the reaction of hydrogen peroxide to water and oxygen gas, including many metal oxides such as manganese dioxide, but the efficiency of catalysts varies. One of the most efficient catalysts for hydrogen peroxide is the enzyme catalase that is contained in blood. (Catalase functions in the body as a peroxide scavenger, destroying peroxide molecules that would otherwise damage cells.)
One catalase molecule can catalyze the reaction of millions of hydrogen peroxide molecules per second. Immediately after each pair of hydrogen peroxide molecules reacts, catalyzed by the catalase molecule, that catalase molecule is released unchanged and becomes available to catalyze the reaction of another pair of hydrogen peroxide molecules. When all of the hydrogen peroxide has reacted to form water and oxygen gas, you end up with as many catalase molecules remaining as you started with.
In this lab session, we’ll measure the reaction rate of the catalyzed reaction of hydrogen peroxide by adding a fixed amount of catalase enzyme to measured samples of hydrogen peroxide. After allowing the reaction to continue for measured periods of time, we’ll stop the reaction by adding sulfuric acid to denature (deactivate) the catalase and then titrate the resulting solutions with a dilute solution of potassium permanganate to determine how much unreacted hydrogen peroxide remains in each sample. In acidic solution, the intensely purple permanganate (MnO4–) ion reacts with hydrogen peroxide to form the light-brown Mn2+ ion according to the following equation:
5 H2O2(aq) + 2 MnO4–(aq) + 6 H+(aq) → 2 Mn2+(aq) + 8 H2O(l) + 5 O2(g)
Because the purple color of the permanganate ion is so intense, the titrant can serve as its own indicator for this titration. As long as hydrogen peroxide remains in excess, MnO4– ions are quickly reduced to Mn2+ ions, and the solution remains a light brown color. As soon as permanganate ions are slightly in excess, the solution assumes a purple color. By determining the amount of permanganate titrant required, we can calculate the amount of hydrogen peroxide that remained in the original samples.
Sulfuric acid is corrosive. Hydrogen peroxide is a strong oxidizer and bleach. Potassium permanganate is a strong oxidizer and stains skin and clothing. Wear splash goggles, gloves, and protective clothing at all times.
If you have not already done so, put on your splash goggles, gloves, and protective clothing.
Transfer about 40 mL of the sulfuric acid to the 100 mL graduated cylinder.
Fill the 10 mL graduated cylinder with the catalase solution.
Label six beakers or other containers from A through F.
Use the pipette to transfer as closely as possible 10.00 mL of the hydrogen peroxide solution to each of the beakers. Record the amount of hydrogen peroxide in each beaker to 0.01 mL on the corresponding line in Table 12-4.
Beaker A is the control, to which we will not add any catalase enzyme. Set it aside for now.
Use the Beral pipette to withdraw 1.0 mL of catalase solution from the 10 mL graduated cylinder.
As close to simultaneously as possible, start the stopwatch or timer and squirt the catalase solution into beaker B. Swirl the beaker to mix the solutions.
With the graduated cylinder of sulfuric acid held ready, when the timer reaches the 15.0 second mark, dump the sulfuric acid quickly into beaker B and swirl to stop the reaction. Record the elapsed reaction time as closely as possible on the corresponding line in Table 12-4.
Refill the 100 mL graduated cylinder with 40 mL of sulfuric acid solution.
Repeat steps 7 through 10 for beakers C, D, E, and F using reaction times of 30 seconds, 60 seconds, 120 seconds, and 240 seconds.
Set up your burette, rinse it with the 0.1 M potassium permanganate titrant, and refill it to near the 0.00 mL line.
Titrate the solution in beaker A by adding titrant to the beaker until a slight purple coloration just persists. Record the volume of titrant required to 0.01 mL on the corresponding line in Table 12-4. (Assuming nominal concentrations, about 35 mL of titrant should be required for beaker A, and correspondingly less for beakers B, C, D, E, and F.)
Repeat step 13 for beakers B, C, D, and E. For each titration, calculate the number of millimoles (mM) of potassium permanganate required to neutralize the remaining hydrogen peroxide. (One millimole is 0.001 mole. Using mM locates the decimal point more conveniently for calculations.) Enter the number of millimoles of titrant required for each titration in Table 12-4.
Using the balanced equation provided in the introduction, calculate the number of millimoles of unreacted hydrogen peroxide in each beaker, and enter that value in Table 12-4.
For each beaker, calculate the reaction rate in millimoles/second (mM/s) and enter that value in Table 12-4.
Hydrogen peroxide | Reaction time | Titrant volume | Titrant millimoles | Peroxide remaining | Reaction rate | |
A. | ________.____ mL | 0.0 s | ________.____ mL | ________.____ mM | ________.____ mM | n/a |
B. | ________.____ mL | ________.___ s | ________.____ mL | ________.____ mM | ________.____ mM | ________.____ mM/s |
C. | ________.____ mL | ________.___ s | ________.____ mL | ________.____ mM | ________.____ mM | ________.____ mM/s |
D. | ________.____ mL | ________.___ s | ________.____ mL | ________.____ mM | ________.____ mM | ________.____ mM/s |
E. | ________.____ mL | ________.___ s | ________.____ mL | ________.____ mM | ________.____ mM | ________.____ mM/s |
F. | ________.____ mL | ________.___ s | ________.____ mL | ________.____ mM | ________.____ mM | ________.____ mM/s |
Q: | Q1: What effect did you observe the catalase catalyst to have on reaction rate? ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ |
Q: | Q2: Based on the data you recorded in Table 12-4, is the effect of the catalyst on reaction rate linear? If not, propose an explanation. ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ |
Q: | Q3: Would you expect the reaction rate to increase, decrease, or remain the same if you increased the amount of catalyst? Why? ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ |
Q: | Q4: When you begin a titration of a reacted hydrogen peroxide solution, you find that the first drop of potassium permanganate titrant causes the solution to assume a purple color. What has happened? ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ |