Rates+of+Reaction+Project

As most of us aren't on MSN at same time, i made this shoutbox. My msn is also blocked, so this is better for me. Just log in and write. (Stole it from Ushud)
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PowerPoint Link
http://www.megaupload.com/?d=FYURG2GZ

Introduction
The rate of reaction is the speed at which a reaction takes place. If a reaction takes more time (is slower) it has a lower rate of reaction. If a reaction takes less time (is faster) the reaction rate is higher. The rate of reaction is calculated by the following equation:

Rate of Reaction= (change in concentration)/time

The four factors that influence the rate of reaction are surface area, temperature, concentration and catalysts.

Surface Area refers to the total exposed area of a reactant. In order to correctly measure the effect that surface area has on a reaction, the mass of the reactant must remain the same (not the volume). As surface area increases (but mass remains constant) the reaction rate will also increase. This is called a direct relationship. This is because with a greater surface area, more of the reactant will come into contact with the other reactant, and therefore more molecules can interact at once. This leads to a greater reaction rate.

Temperature also has a direct relationship with reaction rate. When there is a higher temperature, particles move more quickly. This causes more molecules to come in contact with each other, and therefore react, in a smaller amount of time. Therefore, the reaction rate is greater.

Concentration refers to the ratio of solute verses solvent (such as water, or nonreactive substance) in a solution. The solute is the reactant. When there is greater concentration (more solute vs. less solvent), more of the reactant molecules can come into contact with other reactant molecules (as oppose to solvent molecules) and therefore react more. This means there is a greater reaction rate. This is again, a direct relationship.

A catalyst is a substance that affects the rate of reaction without participating in the reaction. This means that after the reaction has taken place, the catalyst is still chemically unchanged (although it could have undergone a change of state, which is a physical change). Two types of catalysts exist: positive catalysts, and negative catalysts. Positive catalysts increase the rate of reaction, whereas negative catalysts decrease the rate of reaction.

**Purpose** The purpose of this experiment is to investigate the ways in which surface area, temperature, concentration and catalysts affect the rate of reaction. __ Test#1: Surface Area __ The purpose for test is to determine whether a difference in surface will change the rate of reaction. __ Test #2: Temperature __ The purpose of the experiment is to determine if a change in temperature will affect the amount of time taken for a reaction to take place and complete itself. __ Test #3: Concentration __ The purpose for test number 3, was to test whether a change in concentration of one of the __ Test#4: Catalyst __ The purpose of this test is to discover the effect on the rate of reaction when manganese dioxide is added to hydrogen peroxide at room temperature.

**Hypothesis** __ Test #1: Surface Area __ It is hypothesized that when the surface are of the chalk is increased, by changing the chalk from solid chunk to powder, the reaction rate will also increase. __ Test #2: Temperature __ It is hypothesized that when the hydrochloric acid and disodium thiosulphate are mixed at higher temperatures, the rate of reaction will increase, and when mixed at lower temperatures, the rate of reaction will decrease. __ Test #3: Concentration __ It is hypothesized that when the concentration will be increased, so will the reaction time. __ Test#4: Catalyst __ It is hypothesized that when manganese dioxide is added to hydrogen peroxide, the manganese dioxide will prove to be a catalyst and **the rate of reaction will increase**.

Materials
__Test #1: Surface Area__ 3 x Test tubes Scupula 0.5g Solid Calcium Carbonate (Chalk) 0.5g Powdered Calcium Carbonate (Crushed Chalk) 10mL 1M HCl Electric Scale

__Test #2: Temperature__ Hot Plate 5 x 200 mL beakers Scupula Thermometer 5 x Ice Cubes 0.5M HCl - 100mL Distilled Water - 100mL Disodium Thiosulphate - 12.4g 3 x 200 mL of Water Thermometer 3 x Stirring Rods 6 x Test Tubes 3 x Graduated Cylinder

__Test #3: Concentration__

3 test tubes 1.5M HCl - 5mL 3M HCl - 5mL 6M HCl - 5mL 3 x 44mm^2 pieces of Magnesium

__Test #4: Catalysts__

2 125 mL Erlenmeyer flasks 2 50 mL graduated cylinders 2 rubber stoppers with one hole 2 flexible plastic tubes 100 mL water 60 drops 15% hydrogen peroxide solution small amount of manganese dioxide

Procedure
__Test #1: Surface Area__

1) Place 2g each of solid and powered Calcium Carbonate into two separate test tubes. 2) Add 5mL of 1M Hydrochloric Acid into each of the test tubes. 3) Record time for Calcium Carbonate to dissolve.

__Test #2: Temperature__

1) Pour 100mL of water into each of the three 200mL beakers. 2) In one beaker, add 3 ice cubes; leave the other at room temperature; heat the third to approximately 24 oC. 3) Place an empty beaker over a blank sheet of paper. 4) Place a test tube with 20mL of Sodium Thiosulfate into each of the beakers. 5) Place a test tube with 20mL of Hydrochloric acid into each of the beakers. 6) Place a thermometer into the ice beaker to check if it has reached the desired temperature. 7) Repeat for heated beaker. 8) When the temperature reaches 8oC and 24oC respectively, mixed the two chemicals in the test tubes. 9) Record the time it takes until the paper beneath the beakers can no longer be seen through the solution.

__Test #3: Concentration__

1) Add 5mL each of the 1.5M, 3M, and 6M concentrations of HCl into three separate test tubes. 2) Place one piece of Magnesium into each of the test tubes. 3) Record the time it takes for the magnesium to completely dissolve.

__Test #4: Catalyst__

1) In a tub of water or a sink, fill the graduated cylinders with water and suspend them upside down, ensuring there is no air in the cylinders. 2) Place one end of the flexible tubing into the Erlenmeyer flask's rubber stoppers, and the other end into the suspended test tubes. 3) Place 50 mL of water and 30 drops of 15% hydrogen peroxide solution into each Erlenmeyer flask. 4) Add a small amount of manganese dioxide to one flask. 5) Immediately plug both test tubes with the rubber stoppers. Gas produced by the decomposition of the hydrogen peroxide solution should go through the flexible tubing and collect in the inverted test tubes. 6) Record the amount of oxygen gas produced by the reaction with catalyst and the control (without catalyst) over time. 7) At the end of the experiment, use a burning splint to test for oxygen in the flasks and graduated cylinders.

Safety Precautions

 * Throughout the experiments goggles should be worn at all times;
 * Hair must be tied back and sleeves must be rolled up;
 * The area in which the experiment is being performed must be clear;
 * As Test #2 involves a hot plate, necessary precautions must be taken;
 * Care should be taken when handling HCL, as it is corrosive;
 * HCL should not be smelled, as it can damage respiratory organs

Observations
__Test #1: Surface Area__   Table 1
 * Chalk State || Powdered || Solid ||
 * Time for reaction to complete || 1:27:90 min || 16:20:45 min ||

__Test #2: Temperature__

  Table 2
 * || Cold Solution || Room Temperature Solution || Warm Solution ||
 * Temperature (C°) || 8 || 16 || 24 ||
 * Time taken to produce precipitate (s) || 26.43 || 14.6 || 11.09 ||

__Test #3: Concentration__

  Table 3
 * Concentration: || 1.5M || 3.0M || 6.0M ||
 * Time taken for 0.01g of Mg to dissolve: || 9 min. || 1 min. || 20 sec. ||

//~Note: The reason the 1.5M concentration took so long could be due to the fact that the Magnesium floated on the surface and the concentration of an acid tends to decrease at the surface. This may also apply to 3.0M.~//

__Test #4: Catalyst__

(min) ||||  Manganese Dioxide  ||||   Control   || Table 4 After 2 h 45 min, the control sample had produced 14 mm of gas.
 * Time
 * || Amount of gas produced (mL) || Temperature (ºC) || Amount of gas produced (mL) || Temperature (ºC) ||
 * 0 || 0 || 20º || 0 || 19º ||
 * .5 || 3 || 20º || 0 || 19º ||
 * 1 || 7 || 20º || 0 || 19º ||
 * 1.5 || 9 || 20º || 0 || 19º ||
 * 2 || 12 || 20º || 0 || 19º ||
 * 2.5 || 14 || 20º || 0 || 19º ||
 * 3 || 15 || 20º || 0 || 19º ||
 * 3.5 || 17 || 20º || 0 || 19º ||
 * 4 || 18 || 20º || 0 || 20º ||
 * 4.5 || 19 || 21º || 0 || 20º ||
 * 5 || 20 || 21º || 0 || 20º ||
 * 5.5 || 22 || 21º || 0 || 20º ||
 * 6 || 23 || 21º || 0 || 20º ||
 * 6.5 || 24 || 21º || 0 || 20º ||
 * 7 || 24.5 || 21º || 0 || 20º ||
 * 7.5 || 25 || 21º || 0 || 20º ||
 * 8 || 25.5 || 21º || 0 || 20º ||
 * 8.5 || 26 || 21.5º || 0 || 20º ||
 * 9 || 27.5 || 21.5º || 0 || 20º ||
 * 9.5 || 28 || 21.5º || 0 || 20º ||
 * 10 || 28 || 21.5º || 0 || 21º ||
 * 10.5 || 28.5 || 21.5º || 0 || 21º ||
 * 11 || 29 || 21.5º || 0 || 21º ||
 * 11.5 || 30 || 21.5º || 0 || 21º ||
 * 12 || 30 || 21.5º || 0 || 21º ||
 * 12.5 || 30.5 || 21.5º || 0 || 21º ||
 * 13 || 31 || 21.5º || 0 || 21º ||
 * 13.5 || 31.5 || 21.5º || 0 || 21º ||
 * 14 || 32 || 21.5º || 0 || 21º ||
 * 14.5 || 32 || 22º || 0 || 21º ||
 * 15 || 32.5 || 22º || 0 || 21º ||
 * 15.5 || 33 || 22º || 0 || 21º ||
 * 16 || 33.5 || 22º || 0 || 21º ||
 * 16.5 || 33.5 || 22º || 0 || 21º ||
 * 17 || 34 || 22º || 0 || 21º ||
 * 17.5 || 35 || 22º || 0 || 21º ||
 * 18 || 35 || 22º || 0 || 21º ||
 * 18.5 || 35.5 || 22º || 0 || 21º ||
 * 19 || 36 || 22º || 0 || 21º ||
 * 19.5 || 36 || 22º || 0 || 21º ||
 * 20 || 36.5 || 22º || 0 || 21º ||
 * 20.5 || 36.5 || 22º || 0 || 21º ||
 * 21 || 36.5 || 22º || 0 || 21º ||
 * 21.5 || 37 || 22º || 0 || 21.5º ||
 * 22 || 37 || 22º || 0 || 21.5º ||
 * 22.5 || 37 || 22º || 0 || 21.5º ||
 * 23 || 37.5 || 22º || 0 || 21.5º ||
 * 23.5 || 38 || 22º || 0 || 21.5º ||
 * 24 || 38 || 22º || 0 || 21.5º ||
 * 24.5 || 38 || 22º || 0 || 21.5º ||
 * 25 || 38.5 || 22º || 0 || 21.5º ||
 * 25.5 || 38.5 || 22º || 0 || 21.5º ||
 * 26 || 39 || 22º || 0 || 21.5º ||
 * 26.5 || 39 || 22º || 0 || 21.5º ||
 * 27 || 39 || 22º || 0 || 21.5º ||
 * 27.5 || 39.5 || 22º || 0 || 21.5º ||
 * 28 || 39.5 || 22º || 0 || 21.5º ||
 * 28.5 || 39.5 || 22º || 0 || 21.5º ||
 * 29 || 40 || 22º || 0 || 21.5º ||
 * 29.5 || 40 || 22º || 0 || 21.5º ||
 * 30 || 40 || 22º || 0 || 21.5º ||

The burning splint test did not confirm the presence of oxygen when the inside of the Erlenmeyer flask was tested in both the control and manganese dioxide samples.

The burning splint test did not confirm the presence of oxygen when the inside of the test tube was tested in both the control and manganese dioxide samples.

Calculations
__Test #1: Surface Area__ //~No calculations to be amde as no change in concentration~//

__Test #2: Temperature__ //~No calculations to be made.~//

__Test #3: Concentration__ In order to to calculate the rate of reaction we will need to use the integrated rate law:
 * ln([A]0) - ln([A]) = kt ** where ** [A]0 ** is the initial concentration, and A is the final concentration.

We will first need to find the amount in moles of HCl in the first concentration of 1.5M To do this we will use the concentration formula of

We will substitute 1.5 for molarity of solution and .005 for liters of solution and we will call x the moles of solute. In our case x=0.0075 moles of HCl From the reaction equation we will find that Magnesium is the limiting reagent. The 0.01g of Magnesium are equal to 4.1143797572516*10^(-4) moles. This amount in moles will react with 8.2287595*10^(-4) moles of HCl. Subtracting this amount from the original amount of HCl will result in 6.67712404*10^(-3) moles of HCl left unreacted. The molarity of the solution with this much HCl in it would be equal to 1.335M, by applying the formula above. By using the integrated rate law and substituting1.5 for [A]0 and 1.335 for **[A]** and (9/60) for **t**(as t needs to be in hours), we will find the rate, k, to be equal to 0.77477*hour^(-1).

The half life for this reaction would be of t1/2 = ln(2)/k =0.8946 hours

As the experiments are not very accurate, we will find k, whichc should be a constant if only the concentration is changed for all tests.

We will use the same calculations for the other concentrations of HCl, obtaining: k=3.38523hours^(-1) and t1/2 =.204756 hours for the 3M concentration and =5.00623hours^(-1) and t1/2 =0.138456 hours for the 6M concentration.

__Test #4: Catalyst__ average rate of reaction = ∆ concentration/∆ time

The graph to the right shows the concentration of oxygen as a function of time in the solution containing manganese dioxide. Slope is calculated by the following equation: m =∆y/∆x= ∆ concentration of oxygen / ∆ time As gas produced is equal to the concentration of oxygen, the rate of reaction is equal to the slope of the curve in the graph. Thus, the rate of reaction with respect to oxygen produced is equal to the derivative of the function.

The concentration of oxygen as a function of time should be found. Below is an example of how the concentration should be found. First, the results of the amount of oxygen produced in mL must be converted to moles. :

Concentration of Oxygen after 3s: 3 mL n = total volume at STP/GMV n = 0.003 L / 22.4 mol/L = 0.00125 mol

Now the concentration can be found:

M = n/v = 0.00125g/mol/0.003 L = 0.416667 M (this value will be consistent throughout because it is a constant 1 g/22.4 L)

The rate constant can now be found using the integrated rate law. It has a zero order so the equation is as follows:

[A] = [A]0 -kt, where [A] = final concentration, [A]0 = initial concentration, t = time, k = rate constant From t=0.5 to 30

0.416667 = 0.416667 - k(30) 0/-30 = k k=0

These calculations imply that [A]= [A]0. This would mean that the initial concentration is equal to all other concentrations; and basede on the above calculations, this is true as the concentration will be constant throughout.

Below is a chart containing the answers to the above calculations for all of the data.
 * Time (min || Gas Produced (mL) || Moles of oxygen produced || Concentration ||
 * 0 || 0 || 0 || 0 ||
 * 0.5 || 3 || 0.00125 || 1000 ||
 * 1 || 7 || 0.002917 || 1000 ||
 * 1.5 || 9 || 0.00375 || 1000 ||
 * 2 || 12 || 0.005 || 1000 ||
 * 2.5 || 14 || 0.005833 || 1000 ||
 * 3 || 15 || 0.00625 || 1000 ||
 * 3.5 || 17 || 0.007083 || 1000 ||
 * 4 || 18 || 0.0075 || 1000 ||
 * 4.5 || 19 || 0.007917 || 1000 ||
 * 5 || 20 || 0.008333 || 1000 ||
 * 5.5 || 22 || 0.009167 || 1000 ||
 * 6 || 23 || 0.009583 || 1000 ||
 * 6.5 || 24 || 0.01 || 1000 ||
 * 7 || 24.5 || 0.010208 || 1000 ||
 * 7.5 || 25 || 0.010417 || 1000 ||
 * 8 || 25.5 || 0.010625 || 1000 ||
 * 8.5 || 26 || 0.010833 || 1000 ||
 * 9 || 27.5 || 0.011458 || 1000 ||
 * 9.5 || 28 || 0.011667 || 1000 ||
 * 10 || 28 || 0.011667 || 1000 ||
 * 10.5 || 28.5 || 0.011875 || 1000 ||
 * 11 || 29 || 0.012083 || 1000 ||
 * 11.5 || 30 || 0.0125 || 1000 ||
 * 12 || 30 || 0.0125 || 1000 ||
 * 12.5 || 30.5 || 0.012708 || 1000 ||
 * 13 || 31 || 0.012917 || 1000 ||
 * 13.5 || 31.5 || 0.013125 || 1000 ||
 * 14 || 32 || 0.013333 || 1000 ||
 * 14.5 || 32 || 0.013333 || 1000 ||
 * 15 || 32.5 || 0.013542 || 1000 ||
 * 15.5 || 33 || 0.01375 || 1000 ||
 * 16 || 33.5 || 0.013958 || 1000 ||
 * 16.5 || 33.5 || 0.013958 || 1000 ||
 * 17 || 34 || 0.014167 || 1000 ||
 * 17.5 || 35 || 0.014583 || 1000 ||
 * 18 || 35 || 0.014583 || 1000 ||
 * 18.5 || 35.5 || 0.014792 || 1000 ||
 * 19 || 36 || 0.015 || 1000 ||
 * 19.5 || 36 || 0.015 || 1000 ||
 * 20 || 36.5 || 0.015208 || 1000 ||
 * 20.5 || 36.5 || 0.015208 || 1000 ||
 * 21 || 36.5 || 0.015208 || 1000 ||
 * 21.5 || 37 || 0.015417 || 1000 ||
 * 22 || 37 || 0.015417 || 1000 ||
 * 22.5 || 37 || 0.015417 || 1000 ||
 * 23 || 37.5 || 0.015625 || 1000 ||
 * 23.5 || 38 || 0.015833 || 1000 ||
 * 24 || 38 || 0.015833 || 1000 ||
 * 24.5 || 38 || 0.015833 || 1000 ||
 * 25 || 38.5 || 0.016042 || 1000 ||
 * 25.5 || 38.5 || 0.016042 || 1000 ||
 * 26 || 39 || 0.01625 || 1000 ||
 * 26.5 || 39 || 0.01625 || 1000 ||
 * 27 || 39 || 0.01625 || 1000 ||
 * 27.5 || 39.5 || 0.016458 || 1000 ||
 * 28 || 39.5 || 0.016458 || 1000 ||
 * 28.5 || 39.5 || 0.016458 || 1000 ||
 * 29 || 40 || 0.016667 || 1000 ||
 * 29.5 || 40 || 0.016667 || 1000 ||
 * 30 || 40 || 0.016667 || 1000 ||

The balanced chemical equation of the decomposition reaction taking place in this experiment is: 2 H2O2(aq) à 2 H20(l) + O2(g) The rate of reaction in relation to all three of the substances can be expressed as following: Rate = d[O2]/dt = ½ d[H20]/dt = ½ d[H2O2]/dt d[H2O2] = 2 d[O2] = 0.833334 M The average rate of reaction for the control solution is: rate = ∆ concentration/∆ time = 14-0mL/165-0min = 0.08484 mL/min

The average rate of reaction for the solution with the catalyst is: rate = ∆ concentration/∆ time = 40-0 mL /30-0 min = 1.333 mL/min

Conclusion
This lab's main purpose was to establish what changes will also change the reaction rate of the reactants. The changes that were tried were Surface Area Change, Temperature Change, Concentration change and adding a Catalyst. Before the experiment we assumed that the reaction will happen faster if: The results of our experiment proved that our assumption were all correct.
 * the surface area will be increased;
 * the temperature will be increased;
 * the concentration will be increased;
 * a catalyst will be added.

__ Test #1: Surface Area __ In the first test (Surface Area), the reaction rate was increased by changing the chalk from the solid chunk state to the powdered form. The reason for this is that in the chunk state, the hydrochloric acid will not be able to react with the atoms in the middle of the chalk. By breaking the chalk into smaller pieces (powder), there will be a lot fewer atoms that will not be able to react with HCl in the beginning, but they will be able to do so, when the outer atoms of the powder are done reacting with the acid.

__ Test #2: Temperature __ In the second test (temperature), the rate was increased by increasing the temperature of HCl from 8°C, to the room temperature of 16°C and finally to 24°C. The reason for this is that by increasing the temperature, we are increasing the collision frequency. The two reactants will react only if they collide with enough power, and the right way. The minimum energy that the particles need to react is also called the activation energy.

__ Test #3: Concentration __ The reason for the findings of from the third test (concentration) are very simple. In order for a reaction to occur, we will first need to have the two reactants collide. We fulfill this requirement by putting one of the into the other (Magnesium in HCL in our case). As we increase the concentration of HCL, the chances that a collision will happen will also increase and thus the reaction will happen faster. We can also see that the constant "k" is not a constant if the concentration was changed, but that may be becaouse the experiment may not be incredibly accurate.

__ Test#4: Catalyst __ As was hypothesized the manganese dioxide acted as a catalyst and increased the rate of reaction of the hydrogen peroxide. This is very evident from the data. The mixture to which manganese dioxide was added showed an immediate visible reaction. Meanwhile the control mixture showed no visible reaction for at least thirty minutes. In addition, the average rate of reaction of the solution with magnesium dioxide was 0.4849 higher then that of the control solution. If more hydrogen peroxide were to be added to the equation, the concentration would increase. The reason for this relatse to the Maxwell Boltzmann Probability Curve. The reaction is a simple decomposition reaction. Since there is only one reactant, the number of particles of this reactant that are above the activation energy will react, and the particles below the activation energy will not. As the amount of hydrogen peroxide in the solution increases, the percentage of particles above the activation energy will stay the same but the number of them will increase. This also applies in reverse- the less hydrogen peroxide, the lower the reaction rate.

The thermometers were not in the solution during the experiment, which would have produced more accurate results, but were instead measuring the temperature of the air in the Erlemyer Flasks. However there is still an obvious trend- the temperature in both flasks ebgan to rise as the chemical reaction took palce. This is because it is an exothermic reaction, meaning that it produced heat. When teh hydrogen peroxide decomposes, the bonds break and release energy. This energy is measured by the thermometer in the form of heat.

Discussion
__Test #1: Surface Area__  //Other experiment example and real life applications of the surface area findings:// //1.The catalytic decomposition of hydrogen peroxide// This is another familiar lab reaction. Solid manganese(IV) oxide is often used as the catalyst. Oxygen is given off much faster if the catalyst is present as a powder than as the same mass of granules.

//2.Catalytic converters// Catalytic converters use metals like platinum, palladium and rhodium to convert poisonous compounds in vehicle exhausts into less harmful things. For example, a reaction which removes both carbon monoxide and an oxide of nitrogen is:

Because the exhaust gases are only in contact with the catalyst for a very short time, the reactions have to be very fast. The extremely expensive metals used as the catalyst are coated as a very thin layer onto a ceramic honeycomb structure to maximise the surface area.  Explanation of the experiment  The reaction will only result if particles from both reactants will collide with one another. By increasing the surface area, we are also increasing the amount of collides that can happen. As an example we will take Magnesium and Hydrogen gas: __Test #2: Temperature__

**Increasing the collision frequency** Substances react only when their particles collide. By heating the substances, the particles gain more energy and move faster. The frequency of the particle collisions in **gases** is proportional to the square root of the percentage increase in teh Kelvin temperature. For example, if the temperature of the substances were increased from 293 K to 303 K (roughly 20°C to 30&degC), the collision factor of the substances will increase by a frequency of:



The increased rate of collision is 1.7%. This however is not the rate increase of reaction. In fact, in that increase of 10 degrees, most likely the reaction rate has doubled, or increased by 100%. For most reactions that take place around room temperature, an increase of 10 degrees Celcius usually means a doubling of the rate of reaction.

In this experiment, disodium thiosulphate was mixed with hydrocloric acid. In this reaction, when the hydocloric acid is added to the disodium thiosulphate, a yellow precipitate of sulphur is formed. When the two chemicals are completely consumed by the reaction, the sulphur produced is a solid opaque yellow colour.

The only time a reaction will occur from a collision is if there is enough energy in the collision to start the reaction. This minimum energy requirement is known as the **activation energy** for the reaction.

You can mark the position of activation energy on a Maxwell-Boltzmann distribution to get a diagram like this:

Only the particles that have been represented in the green area to the right of the activation energy line will react when collision occurs. The majority of the particles do not have enough energy to start a reaction when collision occurs, so the particles will simply bounce off each other.

To increase the rate of the reaction, there must be a increase in the energy of the particles colliding. By increasing the temperature, the particles now have more energy, and there are then a greater number of particles that have energy levels equal to or greater than the required activation energy.

In the next diagram, the graph labelled T is at the original temperature. The graph labelled T+t is at a higher temperature.



Though there isn't a significantly large movement in the curve, it can be seen that there is now a significant increase (percentage wise) in the number of collisions that have enough energy to result in a reaction. That is to say, there is a significant increase in the area of the curve to the right of the activation energy line.



As stated earlier, the increase in 10 degrees Celcius only results in a 1.7% increase in the number of collisions. That is referring to the total number of **collisions**. Examining the area under the curve that is past the activation energy, it can be seen that there is a new area under the curve **T+t** that is approximately twice the size of the area under the curve **T**. Therefore, it can be concluded that an increase in 10 degrees Celcius can result in approximately a 100% increase in the **rate of reaction**.

__Test #3: Concentration__  //1.The catalytic decomposition of hydrogen peroxide//** Solid manganese(IV) oxide is often used as a catalyst in this reaction. Oxygen is given off much faster if the hydrogen peroxide is concentrated than if it is dilute. //2.The reaction between sodium thiosulphate solution and hydrochloric acid// When a dilute acid is added to sodium thiosulphate solution, a pale yellow precipitate of sulphur is formed. As the sodium thiosulphate solution is diluted more and more, the precipitate takes longer and longer to form.
 * //Other experiment example and real life applications of the surface area findings://

We suppose that we are using a small amount of a solid catalyst in a reaction, and a high enough concentration of reactant in solution so that the catalyst surface was totally cluttered up with reacting particles. Increasing the concentration of the solution even more can't have any effect because the catalyst is already working at its maximum capacity.
 * //There are cases where changing the concentration will have no effect on the reaction rate://**
 * <span style="font-family: Helvetica,Arial;">

<span style="color: rgb(0, 0, 0);"><span style="font-family: Helvetica,Arial;"> __Test #4: Catalyst__

A catalyst is a substance that will affect the rate of reaction in a chemical reaction but will not react itself. When the reaction is complete, the catalyst will remain in its original quantity and composition.

In this experiment, manganese dioxide is a heterogeneous catalyst, meaning that it is a different phase then the reactants. The reactants, water and hydrogen peroxide, are in an aqueous solution. Meanwhile the manganese dioxide is in powder form, and is non-soluble. It will remain a solid once added to the solution.

The reason that manganese dioxide speeds up the rate of reaction is that it weakens the connection between the hydrogen and the oxygen in hydrogen peroxide. When the force bonding the molecules together becomes weaker, it takes less energy for the molecules to separate. This effect is described by the Maxwell Boltzmann Probability Curve. __Maxwell Boltzmann Probability Curve__ <span style="font-family: Georgia,serif;">Graph 1 (source: http://www.chemguide.co.uk/physical/basicrates/introduction.html#top ) The Maxwell-Boltzman probability curve, shown in <span style="font-family: Georgia,serif;">Graph 1, predicts the distribution of particles of different energy levels in a gas. The graph is only accurate for gases; however the general idea can also be applied to liquids. As can be seen, most of the particles have moderate energy levels, some particles have very high energy levels, and very few particles have low energy levels. <span style="font-family: Georgia,serif;">Graph 2 (source:  http://www.chemguide.co.uk/physical/basicrates/introduction.html#top  ) Activation energy (EA) measures the minimum amount of energy necessary for a reaction to occur. Only particles with an energy level above activation energy will undergo a chemical change (<span style="font-family: Georgia,serif;">Graph 2 ). Once the number of particles at each energy level is known, it’s possible to calculate the number of particles that will have a high enough energy level to undergo the chemical reaction. <span style="font-family: Georgia,serif;">Graph 3  ( source: http://www.chemguide.co.uk/physical/basicrates/catalyst.html#top        ) A positive catalyst weakens the intramolecular bonds of molecules in the surrounding substance. Thus the activation energy will be //lower//, because it takes less energy to break the bonds within the molecules. As illustrated in <span style="font-family: Georgia,serif;">Graph 3, when the activation energy is raised, a higher number of particles will be capable of undergoing the chemical reaction. Thus the rate of reaction will increase. Graph 4 illustrates the decrease in the activation energy necessary to cause a reaction.

<span style="font-family: Georgia,serif;">Graph 4 __Practical Applications__ The harber process is used to create ammonium. Nitrogen and hydrogen are combined with iron, a catalyst. This is an example of a heterogeneous catalyst, as the iron is solid while the other components are gas. The equasion is reversable.
 * Harber Process**



Ammonium from the Harber Processnitric is used to creat nitric acid by oxidizing the ammonium. A platinum-rhodium catalyst is used in this reaction to increase the rate of reaction and save money. The catalysts are in large thin sheets to maximize surface area. This way more of the reacting particles make contact with the catalyst at the same time, therefore the rate of reaction can be further increased. The sheet is heated to about 900 degrees celcius, and conveniently enough the reaction is extremely exothermic, and maintains the high temperature. It is then cooled and exposed to the air, at which point the nitrogen oxide oxidizes and forms nitrogen dioxide. Finally, the nitrogen dioxide, still exposed to air, is also exposed to water. It reacts, bonds with hydrogen and finally forms nitric acid.

Nitric acid can be used in metallurgy because it reacts wiht most metals. It reacts with hydrocloric acid to form Aqua Regia, which can dissolve gold and platinum. It's also used in woodworking, to age pine and maple.

Sources of Error
__Test #1: Surface Area__ The error that might have influenced the experiment the most, is the fact that some chalk did not react with the Hydrochloric acid. The reason is that when HCL was added over the chalk, foam formed and the chalk dust sticked to the bubbles in the foam. This was observed in both parts of this experiment with powder and with solid chalk. In the solid chalk part of the experiment, when pieces were breaking from the chalk, they were sticking to the bubbles provided they were small enough. The effect is somewhat balanced, as it happened in both parts of the experiment.

__Test #2: Temperature__ A possible source of error was that the test tubes with hydrochloric acid and disodium thiosulphate were not evenly heated and cooled. The test tubes were not completely enveloped in water, and so when the liquid was heated, it is possible convective actions took place, in which the heated part of the liquids moved to the top of the test tubes and cooled off as it came in contact with air. As well in the cooled test tubes, only the submerged portions of the test tube were fully cooled. The top of the test tube was exposed to air and remained closer to room temperature.

__Test #3: Concentration__ One error in this experiment happened as the Magnesium ribbon floated to the top in two of the tests. While floating, the top part of the ribbon, might have gone out of the solution, reacting with air. Adding to this issue is the fact that an acid's concentration is lower at the top. Although the answers still follow a logical succession (takes more time for magnesium to react in a solution with a lower concentration), the delta t (difference in temperature between two tests), may not be very accurate.

__Test #4: Catalyst__ A source of error in this experiment was the fact that the thermometers were not placed into the reacting solutions, but were rather left above them. Thus the temperature of the air in the Erlemyer flask was measured rather then the temperature of the liquid. This is less accurate. In addition, the equipment may have been a source of error. As can be seen in the observations, although the two thermometers were less then twenty centimeters apart during the experiment, one registered the original temperature to be twenty degress celcius, and the other nineteen. A further source of experimental error is leaky equipment. If the top of the Erlymeyer flask was not on tightly enough, impurities may have entered teh stream of air. The air that was already in the Earlymeyer flask before the experiment began may have been pushed through to the graduated cylindar before the oxygen produced in the reaction. This would explain the failure of the splint test to recognize the presence of oxygen.