chem 18.1 exp 1

Experiment 1Factors Affecting Reaction Rates

Gutierrez, Abonie Jane T., Lomalio, Mae Florence C., Lu, Justine Arnold L.Group 1, MHAB1, Chem 18.1, Mr. Mario Rosete III

In the experiment, the reaction rates were observed while also observing how they were affected by some factors – nature of the reactants, surface area, concentration of the reactants, temperature and the presence of a catalyst. It has been observed that the rate of reaction may vary depending on the nature of the reactant. Some reactant may be more reactive than other reactants. It was also observed that increasing the surface area, temperature, and concentration would also increase the reaction rate. Some substances known as catalysts can either speed up or slow down the reaction without getting used up in the process.

The main objectives of this experiment are to identify the effects of some factors on the reaction rates, to determine the rate law expression by using the method of initial rates, and to evaluate the value of the activation energy of a reaction.

Keywords: reaction rate, concentration, catalyst, surface area, temperature

Chem 18.1 Factors Affecting Reaction Rates of 7

Under a given set of conditions, each reaction has its own reaction rate. Reaction rates are expressed as the change in the concentration of the reactant or the product per unit time. This may refer to the rate at which the product is formed or the rate at which the reactants are consumed.

There are five factors that affect the rates at which a particular reaction occurs. The first factor is nature of reactants. Chemical reactions involve breaking and formation of bonds. Reactants with strong bonds need higher energy to break the bond and form new compounds; while reactants with weaker bonds need lower energy. Some compounds are also more reactive than others. Generally, reactions between ions in aqueous solutions are fast; reactions between covalent molecules are slower.

In order for a reaction to occur, reactants must collide with each other. Reactants can collide only with the surface of the other reactants. Thus, the greater surface area exposed the greater possibility for effective collisions.

The frequency of effective collisions would also increase if we increase the concentration of the reactants. In effect, increasing the concentration would increase the rate of reaction.

The next factor affecting reaction rates is temperature. Increasing the temperature increases the kinetic energy of the reactants. This makes the reactants more energetic than before. With the new acquired energy, reactant molecules will have energy equal to or greater than the Activation energy. Thus, it can be said that an increase in temperature would also increase the rate at which the reaction proceeds.

Some substances can hasten a reaction without getting used up in the process. These substances are called catalysts (positive). A catalyst speeds up the reaction by providing a less energy-requiring path for reaction.

The findings of this experiment are significant in explaining various concepts in chemical reaction. This can also be helpful in other fields of studies, such as, medicine, pharmacy, agriculture and many more,

Experimental

The first part of the experiment focused on the effects of nature of reactants on the rate of

reaction. Two test tubes were labeled A and B. 3 ml of water were placed in each test tube. Mg ribbon was added to test tube A and a small piece of Na to test tube B.

For the next part of the experiment, the effects of concentration of reactants were tested. A 10-ml beaker was placed on a white paper marked with X. 5 ml of 0.15 M Na2S2O3 was added into the beaker. 1 ml of 3M HCl was added. Then, the reaction from the moment the HCl solution was added to the time the X mark is no longer visible was timed.

The same process was repeated but this time using different concentrations of reactants.

On the first part of the continuation of the experiment in concentration, the concentration of HCl was kept constant.4 ml of 0.15 M Na2S2O3 was added with 1 ml H2O, and was added with 1 ml 3M HCl. 3 ml of 0.15 M Na2S2O3 was added with 2 ml H2O, and 1 ml 3M HCl. 2 ml of 0.15 M Na2S2O3 was added with 3 ml H2O, and 1 ml 3M HCl. 1 ml of 0.15 M Na2S2O3 was added with 4 ml H2O, and 1 ml 3M HCl.

On the next part on the test on concentration, the concentration of Na2S2O3 was kept constant, and HCl concentration was the variable. 5 ml of 0.15 M Na2S2O3 was added with 2.5 ml 3M HCl. 5ml of 0.15 M Na2S2O3 was added with 0.5 ml H2O and 2.0 ml HCl. 5ml of 0.15 M Na2S2O3 was added with 1.0 ml H2O and 1.5 ml HCl. 5ml of 0.15 M Na2S2O3 was added with 1.0 ml H2O and 1.5 ml HCl. 5ml of 0.15 M Na2S2O3 was added with 1.5 ml H2O and 1.0 ml HCl. 5ml of 0.15 M Na2S2O3 was added with 2.0 ml H2O and 0.5 ml HCl.

The next part of the experiment focuses on the effect of temperature on the rate of reaction. Two test tubes were prepared-one with 5 ml 0.15 M Na2S2O3 and one with 7.5 ml H2O and 2 ml 3 M HCl the two test tubes immersed in a water bath containing tap water for 5 minutes. The temperature of the water bath was measured and recorded. The reagents were mixed in a 50 ml beaker placed on top of a white paper marked with X. Reaction time was recorded. The same procedure was repeated but this time at two other temperatures.

The next part of the experiment tests how surface area affects reaction rates. Two identical strips of Mg ribbon were taken. One was cut into smaller strips. The uncut Mg ribbon was placed in a test tube and the shredded in another tube. 5 ml of 3M HCl was added to each test tube. Observations were then recorded.


ln Rate vs. ln [Na2S2O3]

-5.0-4.5-4.0

-3.5-3.0-2.5-2.0-1.5

-1.0-0.50.0

-4.0 -3.0 -2.0 -1.0 0.0

ln [Na2S2O3]

ln R

ate

To test the effect of the presence of a catalyst in a reaction, two test tubes were prepared. 5 ml of 3% H2O2 was placed in both test tubes. About one gram Rochelle salt was added to each test tube. A pinch of CoCl2 was added to one test tube.

Results

Part A. Nature of Reactants

TABLE 1. Relative Reactions of Mg ribbonand Na metal with H2O.

Reactants Visible ResultsA Appearance of bubbles on the

surface of the Mg strip; very slow reaction rate

B Formation of bubbles; evolution of gas; faster reaction rate

Upon the addition of Mg strip to 3mL of H2O, it took quite some time before visible results were observed. There was a formation of bubbles on the surface of the Mg strip. If not slow, the Mg didn’t show any notable reaction at all. Meanwhile, the Na metal reacted faster with water than Mg. There was formation of bubbles, liberation of gas, and at the same time a slight bumping within the surface of the test tube caused by a slight explosion.

Since the reaction between Mg and H2O was slower than that of Na with H2O, therefore Mg has higher activation energy (Ea).

The figure below shows a superimposed reaction profile of both Mg and Na with H2O

FIGURE 1. Reaction diagram of Mg and Na.

Part B. Concentration of Reactants

TABLE 2. Time and Rate of Reaction using

Constant HCl Concentration

[Na2S2

O3][HCl]

ln[Na2S2

O3]

t(s) Rate(1/t M/s)

ln Rate

0.125 M

0.5 M

-2.0794 17.70

0.0565

-2.8735

0.1 M 0.5 M

-2.3026 20.17

0.0496

-3.0038

0.075 M

0.5 M

-2.5903 26.97

0.0371

-3.2941

0.05 M 0.5 M

-2.9957 40.30

0.0248

-3.6970

0.025 M

0.5 M

-3.6889 96.25

0.0104

-4.5660

With the concentration of HCl held constant, the data and observations gathered showed that as [Na2S2O3] decreased, the recorded time for each reaction showed a relative decrease in respective rates.

From the table, ln [Na2S2O3] was plotted against ln Rate taking the form of a straight line with a positive slope (FIGURE 2).

FIGURE 2. ln Rate vs. ln [Na2S2O3]

Doing linear regression with ln Rate vs ln [Na2S2O3], will give us a correlation coefficient value of 0.997 (r = 0.997), which showed that there is a good association between rate and the concentration of Na2S2O3.

TABLE 3. Time and Rate of Reaction usingConstant Na2S2O3 Concentration.


EA

EA

Mg ribbon

Na

ENERGY

Reaction Progress

H Change in enthalpy

ln Rate vs. ln [HCl]

-3.3-3.3-3.2-3.2-3.1-3.1-3.0-3.0-2.9-2.9-2.8-2.8

-2.0 -1.5 -1.0 -0.5 0.0

ln [HCl]

ln R

ate

ln k vs. 1/T (in K)

0.0

0.5

1.0

1.5

2.0

2.5

0.0032 0.0033 0.0034 0.0035

1/T (in K)

ln k

[Na2S2O3

][HCl]

ln[HCl]

t(s) Rate(1/t M/s)

ln Rate

0.1 M 1 M 0 16.15

0.0619

-2.7822

0.1 M 0.8 M

-0.2231

16.65

0.0601

-2.8117

0.1 M 0.6 M

-0.5108

21.23

0.0471

-3.0555

0.1 M 0.4 M

-0.9163

22.62

0.0442

-3.1190

0.1 M 0.2 M

-1.6094

25.28

0.0396

-3.2290

As what can be incurred from the table, as the [HCl] decreased with constant [Na2S2O3], the time increased implying decrease in reaction rates.

A graph of ln [HCl] against ln Rate shows a curve line with a positive slope, which theoretically should have been a straight horizontal line.FIGURE 3. ln Rate vs. ln [HCl]

Part C. Temperature

TABLE 4. Time and Rate of Reaction at Different Temperatures.

Temp (0C)

1/T(K) Time (s)

Rate (1/time)

k ln k

16.50

C3.45x10-

3 K98.45s 0.0101 2.9421 1.07

927.30

C3.33x10-

3 K51.35s 0.0195 5.8510 1.76

6370C 3.22x10-

3 K36.84s 0.0287 8.4188 2.13

5

Table 4 shows the time and rate of reaction of Na2S2O3 and HCl at three different temperatures. The data showed that the rate increases as the temperature increases.

The graph of ln k against 1/T (in K) would give us a straight line with a negative slope.

FIGURE 4. ln k vs. 1/T (in K)

Part D. Surface Area

In the test for the effect of surface area on reaction rates, the following data were observed:

Reactants Visible resultsStrip of Mg Formation of heat;

slowly consumed

Pieces of Mg Formation of heat; immediately consumed

Upon the addition of both Mg strip and Mg pieces, formation of heat was observed – a manifestation of an exothermic reaction. However, it was further observed that the pieces of Mg were consumed faster than the strip of Mg.

Part E. Catalyst

The following data were observed in the test for the effect of catalyst on the rate of reaction:

Reactants Visible ResultsH2O2 + Rochelle salt Solution turned from


colorless to greenH2O2 + Rochelle salt + CoCl2

Solution turned from colorless to green, then yellow, and finally brown

Discussion

Under any given set of conditions, each reaction may progress either slow or fast, which is determined by its characteristic rate. Reaction rate is the measure of the quantity of the products formed, or the amount of reactants reacted or consumed per unit time. These relative changes in amounts are measured in terms of concentration.

There are two theories on reaction rates. According to the Collision Theory, equipped with the activation energy (Ea), the minimum energy required for the reaction to be initiated, molecules involved in a reaction must collide in the highest possible frequency in order for the reaction, if not to proceed to completion, to yield more products. Moreover, the lower the activation energy, the faster the reaction. It is the first factor that affects rate of reaction. In the experiment, Mg ribbon reacted faster than the Na metal did. This suggested that its reaction with water has lower activation energy. The following equations show the reaction for each

Mg(s) + H2O(l) --> MgO(s) + H2(g)2Na + 2H2O(l) --> 2NaOH(aq) + H2(g)

Meanwhile, the Transition State Theory is based on the postulation that in all reactions, before finally ending up with the desired products, an intermediate complex is first formed. The rate of the formation of such transition state is equal to its rate of decomposition, which makes it short-lived and hard to be isolated. The rate of reaction is directly proportional on the rate of the formation of the transition state complex. Therefore, the faster, or the lower energy needed, for the complex to form, the faster the reaction would occur, and vice versa.

Another factor that affects rate is the concentration of reactants. In the experiment, consider the reaction

Na2S2O3(aq) + 2HCl(aq) ---> 2NaCl(aq) + SO2(g) + H2O(l),

which shows that as [Na2S2O3] decreased, with [HCl] held constant, the recorded time for each reaction showed a relative decrease in

respective rates. Doing linear regression with ln Rate vs ln [Na2S2O3], will give us a correlation coefficient value of 0.997 (r = 0.997), which showed that there is a good association between rate and the concentration of Na2S2O3. The order of the reaction with respect to Na2S2O3 is second (b = 2.28 2).

Furthermore, as the [HCl] decreased with constant [Na2S2O3], the recorded time increased. Theoretically, there should not have been change in the rate of reaction since the order with respect to HCl is zero (b = 0.29 ). Ideally, the graph of ln [HCl] vs ln Rate should have been a horizontal line (b = 0). This deviation may be due to some inaccuracies with regards to the amount and concentration of Na2S2O3 delivered in every system. The overall order of the reaction between Na2S2O3 and HCl is two (2 + 0 = 2). Deriving the rate law expression:

Rate = k [Na2S2O3]2 [HCl]0

or simplyRate = Rate = k [Na2S2O3]2

The specific rate constant is

from Rate = k [Na2S2O3]2

k = Rate [Na2S2O3

2] = 0.0619 M/s (0.1 M)2

= 6.19 M-1 s-1

This relationship observed between concentration of reactants and rate is also based on the Collision Theory. The more molecules there are, the more frequently they collide, and the more often or faster a reaction occurs. Thus, reaction rate and concentration of reactants are directly proportional:

Rate collision frequency concentration

We looked at a simple, one step reaction, in which reactant molecules collide and form products, but even the rates of complex reactions depend on reactant concentration.

Molecules must possess enough energy in order for them to react. Temperature as another factor in the rate of a reaction has a major effect on the kinetic energy of the molecules, and thus increasing the energy of and for the collisions. In the jumble of molecules in the reaction of Na2S2O3 with HCl, most collisions occur with so much energy that the


molecules react, and repeating the reaction in a much lower temperature will lower the probability of sufficiently energetic collisions to occur. Thus, lowering the temperature will slow down a reaction. On the other hand, an increase in temperature within the system in which a reaction occurs increases the reaction rate by increasing the number and, especially, the energy of the collisions:

Rate collision energy temperature

To calculate the value of the Ea in the reaction using Arrhenius equation,

K = Ae –Ea/RT

ln k = ln A -Ea/R (1/T)

-4589= -Ea/ 8.3145Ea= 38160 J/mol

Physical state or the nature of the reactants also intensively affects the rate of reaction. Molecules must mix to collide, and in order for them to mix, it is necessary that they be of the same phase. Reactants of the same phase, as in an aqueous solution, contact of molecules is due to thermal motion. However, those in different phases have contacts only at the interface, so there’s a need for vigorous shaking and reduction of size by means of grinding. In the experiment, pieces of Mg reacted with HCl faster than strip of Mg. The more finely divided the reactant is, the larger surface area it has per unit volume giving more sites for collisions, and the faster the reaction occurs.

Catalysts are substances that either speed up or slow down reaction rates without being consumed in the system. In this case for the combination of Rochelle salt/Potassium Sodium tartrate and hydrogen peroxide, the positive catalyst for this experiment is the cobalt (II) chloride.

Without the Catalyst, the experiment should have an equation that goes like this:

5H2O2 (aq) + C4H4O62-

(aq) → 4CO2 (g) + 2OH – (aq)

+ 6H2O (l)

or we can divide it into two parts:

3H2O2 (aq) + C4H4O62–

(aq) → 2CO2 (g) + 2HCOO–

(aq) + 4H2O (l)

2HCOO– (aq) + 2H2O2 (aq) → 2CO2 (g) + 2H2O (l) + 2OH–

(aq)

______________________________________

5H2O2 (aq) + C4H4O62-

(aq) → 4CO2 (g) + 2OH – (aq)

+ 6H2O (l)

In the experiment, the group discovers that the mixture starts at a colorless state then it becomes green for while then reverts back to its original color.

With the addition of CoCl2, the mixture, there was first the formation of three layers. Then after a little bubbling, the solution turned from a pinkish color to a dark green. This is caused by first ionizing the catalyst into Co2-

ions. Then the pink Co2- ions are oxidized by the strong H2O2 into green Co3- ions. It’s given by this equation:H2O2 (aq) + 2H+

(aq) + 2Co2+ (aq) 2H2O (l) + 2Co3+

(aq)

After a while, the green Co3- ions then oxidize the tartrate to revert back to the pinkish Co2- ions. During the reaction, the colors of the two ions mix and form a brown solution until it gradually becomes pinkish brown indicating that the reaction did not proceed to completion.


ConclusionsTo sum it all up, the group has

discovered that reaction rates can be affected by five factors namely nature of reactants, concentration, temperature, surface area, and catalyst. With regards to the nature of the reactants, the lower the activation energy of the reactants, the faster the reaction occurs so the relationship between activation energy and reaction rate is inversely proportional.

A higher concentration of reactants means that there will be more chances of effective collisions to occur and proceed to making products. Therefore, the reaction rate is proportional to the concentration of non-zero-order of reactants used.

Temperature, aside from the measure of heat, is also the measure of kinetic energy. So a higher temperature excites the molecules making them move faster. This also gives the molecules greater energy, a boost to achieve the activation energy. In short, the higher the temperature, the faster the reaction rates.

The greater the surface area, the more exposed two substances are to one another so there is more room for collisions to occur. Particle size is also associated with surface area and is often the basis of tests. The group has found that the surface area is also directly proportional to the reaction rate.

The last item, catalyst, may have two kinds of effects on the reaction namely speeding it or slowing it down. The substance that does the former is called a positive catalyst while the latter is called a negative catalyst. So the relationship between a catalyst and the reaction rate doesn’t really have a standard description.

RecommendationsThe most notable mistake the group

made is that of the experiment of constant [Na2S2O3] with varying [HCl]. The results proved to be far from the theoretical because they didn’t observe proper measurements of the reactants. It should have been done it a much more careful attitude.References:

Silberberg, Martin. Principles of GeneralChemistry. New York: McGraw HillBook Company, Inc., 2007. 674-678.

I hereby certify that I have given substantial contribution to the report.

___________________________Abonie Jane T. Gutierrez

__________________________Mae Florence C. Lomalio

__________________________Justine Arnold L. Lu


chem 18.1 exp 1

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