Gas

Design Lab #17: Reaction Rate – Calcium Carbonate & Hydrochloric acid

Design D: Introduction: In this experiment calcium carbonate will be put into a flask and mixed with hydrochloric acid to produce calcium chloride, water, and carbon dioxide.

The formula for this reaction is:

CaCO3(s) + 2HCl(aq) > CaCl2(aq) + CO2(g) + H2O(l)

Purpose: The purpose of this experiment is to determine how the surface area of CaCO3(s) affects the rate of reaction by measuring the volume of CO2(g) produced with time.

Background: When solid reactants are mixed with liquid reactants only the particles on the surface of the solid will have direct contact or collide, to the other liquid reactant particles. When there is more surface area there will be more solid particles exposed to collide with other liquid particles. When there are small pieces of the solid reactant, the rate would be faster than if it were larger pieces with the same mass. There will be more collisions per unit of time, which means the reaction will proceed faster.

Variables: Independent (Changed) ) Surface area of CaCO3 Dependent (Measured)

1. The volume of the gaseous product formed (CO2(g)) Controlled (Constant). Mass of CaCO3
2. Temperature of reaction
3. Concentration of HCl
4. Volume of HCl
5. Time intervals for recording volume

Apparatus: 100 mL

Conical flank single-holed rubber stopper 90 g of CaCO3 chips90 g of CaCO3 powder 90 g of CaCO3 tablets100 mL gas syringe 100 mL graduated cylinderDigital Stopwatch 450 ml 1. 0 M HCl Stand & Clamp Electronic balance.

Safety Procedures:

1. Wear safety goggles for protection.
2. Handle HCl acid with care. If splashes on skin wash immediately
3. Always point gas syringe downwards.

Procedure:

1) Set up equipment for the experiment. Set up the gas syringe with the clamp and connecting pipe connecting to the flask. Have the rubber stopper and stopwatch nearby at your station.

2) Place an on an electronic balance.

3) Place the CaCO3 tablet on and weigh 10 g precisely on an electronic balance. Record mass.

4) When taking mass make sure the tablet, chips, and powder all have the same mass; 10g

5) Measure 50 mL of acid with a graduated cylinder. Pour into conical flask.

6) For the first reaction use the CaCO3 tablets. Start stopwatch immediately after CaCO3 tablets are added into flask. Simultaneously cover the flask with stopper.

7) Ensure that the connecting pipe from the flask to the syringe is connected properly.

8) At 10 seconds record the volume of gas in the syringe.

9) Record the volume of gas inside the syringe every 10 seconds until you have 3 consistent readings.

10) Repeats steps 2 through 8 for two more trials to have 3 values for every type of CaCO3.

11) Repeats steps 2 through 7 for the CaCO3 chips

12) Repeats steps 2 through 7 for the CaCO3 powder.

Method:

1. Cut magnesium ribbon into 15 20mm strips ±0. 5mm
2. Fill measuring cylinder with 100ml ±0. 1ml water. Invert inside an ice-cream container. Fill container with water.
3. Take the 5 mol dm-3 acid and pipette 5ml ±0. 1ml of acid into each 3 test tubes. Insert the delivery tube into a measuring cylinder underwater and prepare bung above the boiling tube.
4. Pour one test-tube into the boiling tube. Place one piece of 2cm ±0. 5cm magnesium into the tube, affix bung to the tube and begin the stopwatch
5. After 5 seconds, record the amount of water displaced on the measuring cylinder. This will be in milliliters ±0. 1ml. Repeat this at 10 seconds.

Continue recording at 5-second intervals until 3 consistent readings are gained. Repeat steps 4 through 7 with the other two prepared test tubes to give 3 readings per concentration

Repeat steps 3 through 8 with the remaining four prepared concentrations of acid Stat stopwatch Place the marble chips and powdered marble into separate test tubes. Add 10cm of the dilute hydrochloric acid to each of the test tubes and observe the rate at which carbon dioxide is produced Gas Syringe Method Equipment Conical Flask, Bung, Connecting Pipe, Gas Syringe, Hydrochloric Acid,  Magnesium Ribbon, Clamp, Stand. For this method we set up the equipment as follows; The first to do is to gather together all equipment, then once that is done get the conical flask and put the bung into the top of it. Then connect the pipe from the bung to the gas syringe, making sure that it is airtight. Then once everything is set up get the stop clock ready,  measure out the 50ml of hydrochloric acid and pour it into the conical flask and again the same as the burette method simultaneously put the magnesium into the conical flask, put the bung into the top of the flask and start the timer. We took down measurements in 5-second intervals. Does the gas syringe have a volume of 100ml?

Vapor Pressure and Heat of Vaporization

Introduction

Evaporation is the process of a liquid becoming vaporized. When a liquid is placed into a confined space some of the liquids will evaporate. Evaporation of the liquid depends on the strength of the intermolecular forces that are between liquid molecules. During the evaporation process of the liquid, new gas molecules exert pressure in the sealed container, while some of the gas condenses back to the liquid state. If the temperature inside the container is kept constant, then the equilibrium at some point will be reached. When the equilibrium is reached, the rate of condensation is equal to the rate of evaporation and the rate of vapor pressure will remain constant as long as the temperature in the sealed container does not change. ?The relationship between the vapor pressure of a liquid and temperature is described in the Clausius-Clayperon equation: lnP=? HVAC / R (1/T)+C. where 1nP is the natural logarithm of the vapor pressure. Hvap is the change in heat vaporization, R is the universal gas constant, which is (8. 31 J/mol•K), T is the absolute, or Kelvin, temperature, and C is the constant that is not related to heat capacity.

Therefore, the Clausius-Clayperon equation does not only describes how vapor pressure is affected by the temperature but relates to the factors of heat vaporization of a liquid. ?The purpose of this experiment is to determine the relationship between the pressure and temperature of the volatile liquids. The pressure will be measured in a sealed vessel that contains different types of liquids such as methanol, ethanol, and propanol. It will be measured several times at different temperatures. At the conclusion of this experiment, the heat of vaporization will be able to be calculated.

Materials: To be able to complete this lab procedure, the materials that are needed is a Vernier computer interface, a Vernier Gas Pressure Sensor, temperature probe, rubber stopper assembly, plastic tubing with two connectors, hot plate, ice, one twenty milliliter syringe, one 400 milliliter beaker, two 125 milliliter Erlenmeyer flasks, one 1 liter beaker, ethanol, methanol, and 1-propanol.

Methods: The first step in performing this experiment is to obtain and wear goggles. The alcohols used in this experiment are flammable and poisonous. The second step is to obtain the materials that are needed and set them up as accordingly. The third step is to use a hot plate to heat 200 milliliters of water in a 400-milliliter beaker. The fourth step is to prepare a room temperature water bath in a 1-liter beaker. The fifth step is to connect the Gas Pressure Sensor to channel one of the Vernier computer interfaces, then connect the Temperature Probe to channel two of the interface and then connect it to a computer. The sixth step is to use the clear tubing to connect the white stopper to the Gas Pressure Sensor.

The white stopper must be twisted snugly into the neck of the Erlenmeyer flask, to avoid losing any of the gas that will be produced when the liquid starts evaporating. The most important thing to do is to remember to close the valve on the white stopper. ?The seventh step is to draw in 3 milliliters of methanol into the 20-milliliter syringe that is part of the Gas Pressure Sensor accessories. Place the syringe onto the valve of the white stopper. The eighth step is to start the Logger Pro program and open the file “34 Vapor” from the Advanced Chemistry with Vernier folder. The ninth step is to click “collect” to begin collecting data. The first measurement will be the pressure of the air in the flask and the room temperature. Place the Temperature Probe near the flask. When the pressure and temperature readings are stabilized, click “keep” to record the readings. The tenth step is to add methanol to the flask by opening the valve below the syringe, push down on the syringe to inject the 1-propanol, and quickly close the valve. Afterward, remove the syringe from the stopper and monitor the pressure and temperature readings. The eleventh step is to place the stoppered flask into the 1-liter beaker of room temperature water.

Place the Temperature Probe in the water bath and monitor the pressure and temperature readings. The twelfth step is to add a small amount of hot water to warm the water bath by only a few degrees. Stir the water with the temperature probe and monitor the pressure and temperature readings. For the thirteenth step, repeat step twelve until five trials are completed. Add hot water for each trial so the temperature of the water bath increases. After the fifth trail is recorded, open the valve to release the pressure in the flask and dispose of the alcohol as directed. The fifteenth step is to end the data collection and record the pressure and temperature readings in the data table. When recording the data, record the pressure valve of the first data point as Pair for trials one and two and record the temperature for trial one. Record the pressure value of the second data point as Ptotal for trial two as well as the temperature. The remaining values are recorded as Ptotal for trial two as well as the appropriate temperature. The last and final step is to clean the work area.

Data Table:

Discussion At the end of this experiment, the results we obtained varied because of the different temperatures and pressures that we observed. During the evaporation process of the liquid, gas molecules exert pressure in the sealed container, while some of the gas condenses back to the liquid state. If the temperature inside the container is kept constant, then the equilibrium was reached. When the equilibrium is reached, the rate of condensation is equal to the rate of evaporation and the rate of vapor pressure will remain constant as long as the temperature in the sealed container does not change.

1. A sample of oxygen of mass 25. 0 g is confined in a vessel at 0°C and 1000. torr. Then 6. 00 g of hydrogen is pumped into the vessel at constant temperature. What will be the final pressure in the vessel (assuming only mixing with no reaction)?
2. A gaseous mixture contains 3. 23 g of chloroform, CHCl3, and 1. 22 g of methane, CH4. Assuming that both compounds remain as gases, what pressure is exerted by the mixture inside a 50. 0-mL metal container at 275°C? What pressure is contributed by the CHCl3?
3. A study of climbers who reached the summit of Mt. Everest without supplemental oxygen revealed that the partial pressures of O2 and CO2 in their lungs were 35 torrs and 7. 5 torrs, respectively. The barometric pressure at the summit was 253 torr. Assume that the lung gases are saturated with moisture at a body temperature of 37°C. Calculate the partial pressure of inert gas (mostly nitrogen) in the climbers’ lungs.
4. During a collision, automobile airbags are inflated by the N2 gas formed by the explosive decomposition of sodium azide, NaN3. 2NaN3 –> 2Na + 3N2. What mass of sodium azide would be needed to inflate a 25. 0-L bag to a pressure of 1. 40 atm at 25°C?
5. Calculate the volume of methane, CH4, measured at 300. K and 825 torr, that can be produced by the bacterial breakdown of 1. 10 kg of simple sugar. C6H12O6 –> 3CH4 + 3CO2
6. We burn 12. 50 L of ammonia in 20. 00 L of oxygen at 00. °C. What volume of nitric oxide, NO, gas can form? What volume of steam, H2O(g), is formed? Assume that all gases are at the same temperature and pressure and that the limiting reactant is used up. 4NH3 (g) + 5O2 (g) –> 4NO(g) + 6H2O(g)
7. A particular tank can safely hold gas up to a pressure of 44. 3 atm. When the tank contains 38. 1 g of N2 at 25°C, the gas exerts a pressure of 10. 1 atm. What is the highest temperature to which the gas sample can be heated safely?

Finding the Ratio of Moles of Reactants in a Chemical Reaction Purpose: The goal of the lab is to determine the mole ratio of two reactants in a chemical reaction (AgNO3 and K2CrO4). However, the formulas for the products are unknown. Introduction: When determining the molar ratio of a chemical equation, usually the formulas of the reactants and the products are known. With that information, it is particularly easy to determine the ratio. However, since the products and the formulas for the products are unknown, another property of the reaction must be analyzed to find the ratio.

This property depends on the amount of the product formed or on the amount of reactant that remains. Properties may include the color intensity due to the product, the mass of the precipitate that forms, or the volume of a gas evolved. In this experiment, the method of continuous variations will be used to determine the mole ratio of two reactants. With this method, the total number of moles of reactants is kept constant for the series of measurements. The property that is going to be measured is the change in temperature.

The temperature change, or the heat produced, will be directly proportional to the amount of reaction occurs and to the total extent of it. The optimum ratio, which is the ratio of the reactants in the balanced chemical reaction, will form the greatest amount of product, or generate the most heat, and will be key to determining the molar ratio. Corrosive liquids, which burn the skin, will be used in the experiment. When this liquid reacts with acid, a toxic gas will formed. Keep away from the gas and protect your skin and clothing.

Work in a fume hood or well-ventilated lab. Wear chemical splash goggles, chemical-resistant gloves, and a chemical-resistant apron. Wash hands thoroughly with soap and water before leaving the laboratory. The molar ratio of the reactants is the ultimate goal of the lab. In order to achieve that, secondary observations on the temperature change will have to be made and analyzed. The data and graph made after the data is attained will assist in that. Procedure: 1) Obtain 2 beakers with 175mL of NaClO in one and 175mL of “Solution B” in another. ) Measure the temperature of each and make sure they are the same. 3) Measure 5. 0mL of NaClO and 45. 0mL of “Solution B” with the appropriate graduated cylinders and add them to a Styrofoam cup. 4) Stir the solution with a thermometer, and record the max temperature reached. 5) Pour the solution out, rinse the cup, and repeat steps 1-4 using a different ratio of the two substances, keeping the total volume at 50. 0mL. 6) Continue testing various ratios until you have at least 3 measurements on either side of the peak temperature difference.

Conclusion: When the formula of the products are unknown in a chemical reaction, experiments must be done to find the mole ratio of the reactants. In our experiment, we used the method of continous variations to determine to the mole ratio of the two reactants. The property measured was the change of temperature, as indicated in the data table. The method of continous variations keeps the total number of moles of reactants constant through a series of titrations.

Each titration varies the mole fraction of each reaction from mixture to mixture by adjusting the ratio of NaClO to Na2SO3, which is also indicated in the data table. Theoretically, the maximum temperature change occurs when teh mole fraction of the reactants is closest to the actual stoichiometric mole ratio, which signals the mole ratio based on the mole fraction in the titration. According to the analysis, the mole ratio is 1:2 in the order of NaClO to Na2SO3.

This ratio was concluded by the graph, in which the lines of best fit were extrapolated to intersect at the optimum ratio point. However, there was room for error when measuring the liquids in each titration. The measurements weren’t always exact which could affect the change in temperature. Also, when measuring the temperature it might not have been exact due to inconsistent stirring. Nonetheless, the goal of the lab was to find the mole ratio of the two reactants and it was concluded to be 1:2.

I am going to investigate how concentration of hydrochloric acid affects the rate of reaction between hydrochloric acid and indigestion tablets which contain mainly calcium carbonate.

From my background knowledge from class work and books, (see references) I have found out that if you increase the concentration of hydrochloric acid, the rate of reaction will increase and the time of the reaction will decrease. The concentration is dependant on the proportions of hydrochloric acid and water in the solution. The stronger the hydrochloric acid is, the higher the concentration is. I know from my research that other things can affect the rate of reaction, for example:

Temperature of acid- the higher the temperature of the acid is, the more energy the particles have to move around, therefore there are more collisions and so a faster rate of reaction. There is a certain amount of energy needed for the particles to react which is called the activation energy, so when the temperature of the solution is higher, it gives more particles sufficient energy so they move faster to react when they collide more. Size of the particles- when the reactant is a solid then it can be broken down into smaller pieces or into a powder giving it different surface areas.

The smaller the pieces, the bigger the surface area is and therefore there is more area for the acid to react with it, and so there is more chance of the particles colliding, so the rate of reaction will increase. Catalysts- this weakens the bonds in the reacting molecules so it seems to lower the activation energy for the reaction. This means that there can be many more successful collisions because particles will have more energy than the activation energy, and so the reaction will be faster. In order to keep my experiment fair, I must keep all the variables the same except concentration, which is what I am investigating.

From my preliminarty experiments, I have found that a gas is let off in this reaction and having testing by putting it with lime water, I have concluded that the gas is carbon dioxide because the lime water turned cloudy. The equation is: Calcium Carbonate + Hydrochloric acid Calcium Chloride + water + carbon dioxide To find the rate of reaction, I will measure how long it takes to produce a certain amount of gas. To work out the rate of reaction, I have to divide the amount of gas I will collect with the time.

Rate= amount of gas collected/ time In order to make this experiment fair, I will keep all the variables the same, except concentration. I will keep the mass of calcium carbonate the same by using one tablet each time. The masses of each tablet vary, but only by a tenth of a gram either side of 1. 01g, which I think is not a large enough difference in mass to make a difference to my experiment. I will try to do all my experiments on the same day, so the room temperature will be the same, which means the temperature of the acid will not change.

I will use a burette to measure the amount of hydrochloric acid and water, so the volume of acid will be the same. I will keep the surface area the same because I will use the whole tablet and so each one will have the same surface area because they are all similar in size. Finally, I will use the same apparatus throughout my whole experiment to make it a fair test. I have done preliminary experiments in order to find the amounts I should use for the variables. I chose to collect 40cm? of gas, 50cm? of acid solution and use half a intigestion tablet and crush it.

I found three problems with using this half a crushed tablet. The first problem was the fact that it did not react very strongly, and it did not collect more than 24cm? of gas with my lowest concentration of acid, secondly it was difficult to get exactly half a tablet, and this would take too long in my real experiment if I was to get exactly half a tablet each time. Lastly, I did not know when to start my stop clock, because the time delay from the first bits of calcium carbonate falling into the acid, to the last bits of calcium carbonate falling was quite long and in between this, some gas was lost.

This has made me decide to use a whole intigestion tablet, so I do not loose as much gas inbetween putting the tablet into the solution and putting the bung on the conical flask. GRAPH I decided that the lowest concentration I will use is 1M of hydrochloric acid. , which took 85 seconds to collect 40cm? of gas. This highest concentration I will use is 3. 8M of hydrochloric acid, which took 50 seconds to collect 40cm?. From my preliminary experiments I have decided to: Use one whole tablet, 50cm? of different acid concentrations and time how long it takes to collect 40cm? f carbon dioxide gas. My prediction is: The higher the concentration of hydrochloric acid, the quicker the reaction time is with the indigestion tablets. This is because I have found out that the reaction will be quicker as the concentration increases, because the higher the concentration is, the more particles of acid there are which are closer together to collide more with each other and therefore react with each other, and so the reaction will be quicker. This is called the collision theory.

I will use a graph to show my results, and I know from previous knowledge that the graph should be directly proportional, which means as the concentration increases, so does the rate of reaction. If the concentration is doubled, the rate of reaction is doubled, because there is twice as much possibility for collisions because there are twice as much particles. The line of best fit should go through the origin because when there is no concentration of acid, there are no particles to react. I expect to get a graph which looks like the following:

Equipment · 1 burette containing hydrochloric acid · 1 burette containing water · 1 conical flask · bung and deliver tube · gas syringe · stop clock · clamp stand I am using burettes because they have an accuracy of 0. 1cm? which means I can measure the volumes of water and acid very accurately, and the range of the burette is 0-50cm? of liquid, which is enough for what I need. I will use a gas syring which is accurate to 1cm? of gas collected, and ranges from 0-100cm? of gas which is accurate enough if I am to collect 40cm? of gas.

The stop clock is accurate to the 100th of a second, but I will round the time to the nearest second, because it is more realistic when remembering human reaction times, which is about 0. 1 seconds. In order to not let any gas escape, I will make sure I put the delivery tube and gas syring securely together. Method · Set up apparatus as above · Take 50cm? of the following concentrations at one time, using the burettes of acid and water-1M, 1. 4M, 1. 8M, 2. 2M, 2. 6M, 3. 0M, 3. 4M, 3. 8M. · Put the acid solution into a conical flask · Put 1 intigestion tablet into the the acid and put the bung on. Start the stop clock and time until the marker reaches 40cm?. · Do this for all the concentrations. I will use a range of 2. 8M of hydrochloric acid, the lowest concentration is 1M and the highest is 3. 8M. I decided to use these concentrations, because in my preliminary experiments, I saw that the reaction was too slow with a concentration below 1M, and that the reaction would be too fast above 3. 8M. I have chosen to do 8 different concentrations, because I will not have enough time to do more, and I will still beable to draw a concusion even if I only use 8 different concentrations.

I will have to use both 2M and 4M hydrochloric acid in order to make the different concentrations of acid. The ones which are 2M or below I will make with the 2M hydrochloric acid, and for the rest 4M hydrochloric acid. I will try to use as little of the 4M acid as possible, because it is more dangerous than the 2M. I will take as many repeat readings as I can in the time that I have, because repeats will help me to make sure I do not get any anomalous results. I will reapeat the anomalous results first. The more repeats I do, the more reliable my results will be.

Safety I will use 2M and 4M hydrochloric acid which both have IRRITANT warnings so I will be careful using them and try not to get them on my hands or in my eyes. I will use goggles to protect my eyes. GRAPH Results This conclusion supports my prediction well because my results show that the higher the concentration, the quicker the reaction is because there are more particles to react with each other, and so there is more chance for them to collide and therefore the reaction is faster, which is what I originally assumed in my prediction.

My graph shows it is directly proportional, because if I take the concentration of 1. 5M of hydrochloric acid, and find the rate of reaction using my graph, it shows that the rate of reaction is 0. 36 cm? /s, and using the graph if I double the concentration to 3M, the rate of reaction is 0. 72 cm? /s which is exactly two times faster than the reaction with 1. 5M which shows it is directly proportional. In the following table, I have calculated the average time and rate of reaction for all the different concentrations.

I have then worked out the difference between each rate in order to find if there is a trend in how much quicker the reaction is which each concentration. GRAPH From this table, I can see there is a trend, because as the concentration goes up by 0. 4M each time, the rate goes up by 0. 07, 0. 08 or 0. 09 cm? /s which are very close to each other and shows that the rate is quite consistent because no matter what the concentration is, the rate goes up in a certain way on average of 0. 08 cm? /s. The only results that do not go with trend are the 3. M concentrations. On my graph I have circled them as anomalous results. There can be several explanations for this which I will cover in the evaluation. The following diagram is a simple way to help show why the rate of reaction increases with the concentration: My experiment has helped me with my conclusion that the rate of reaction increases as the concentration of the hydrochloric acid increases, and has given me evidence to help explain it. Evaluation My results are as realiable as I could make them using the apparatus and the time I had.

From my results I can say that most of the results are quite reliable and accurate to what they should be because I got the results I expected. However, I did get two results which I would say are anomalous. I decided that these two results are anomalous because according to my background knowledge and the rest of my results, I knew that I should get a directly proportional line of best fit, and the rest of the results are very near to this line of best fit. I know that my line of best fit is correct because as the concentration doubles, the rate doubles.

The results for the 3. 8M showed that the rate was slower than the rate of reaction with a lower concentration of 3. 4M. There are many different factors which may have affected my results. One of the biggest faults in my experiment was the fact that I did not have enough time to complete it in one day. Due to various problems, I had to do the experiment on three different days. This means that all the equipment was different which may mean that they work differently from eachother. This makes it an unfair test.

The second problem with doing it on different days is the problem of room temperature which can have a big effect on the rate, because as I know from back ground knowledge, I know the warmer the acid is, the faster the reaction because particles have more energy so there are more successful collisions. I made the mistake of not recording which results are from which day, so I cannot tell if this had a major effect on the results. Whilst doing the experiment, I noticed a few problems which may also have effected my results.

First of all is the problem that I only have two hands, so it was difficult to put the tablet in the conical flask, close the bung and also start the stop clock, all at the same time. When there was someone available, I asked them to start my stop clock, but this was not possible all the time. Adding this time to human reaction time of around 0. 1 of a second, some time could have been lost. Some gas was also lost in the time period between putting the tablet in and putting the bung on. I tried my best to make this time period very small, but still some gas was lost.

When I had managed to get the tablet into the acid with the bung on and time it, I noticed that sometimes the whole tablet would not go into the acid, and so it was not all reacting, so in order for the whole tablet, I would shake it for a couple of seconds. I did not count how long I would do this for each one, but when I did shake it a lot of gas would be produced, so If I shook one flask for longer, more gas would be produced faster because the whole tablet would be reacting with the acid and there would be more collisions and therefore a quicker reaction.

Between each different concentration, I would wash the conical flask, and I observed that if I washed the flask with hot water, the flask would become hotter, or if I washed it with cold water the opposite would happen. This meant that the temperature of the acid and water solution would vary. This made the tests unfair because if I did some of them with hot conical flasks and others with cold ones, the ones with the warmer flasks would react faster because the temperature of the acid would increase and so give the particles more energy to react.

If the equipment was much more sophisticated, for instance if all the equipment would stay the same temperature or if there was special clock which would start at the exact time the tablet touched the acid, my results would be much more accurate, but I still found good results. [IMAGE]If I could do the experiments again, I would do the following things differently in order for my results to be more accurate.

I would make sure I did them all on the same day, use all the same equipment, have someone to start the stopclock, have better equipment, for instance a conical flask with a divider so the acid and calcium carbonate won’t mix until I want them to: Apart from all of the problems, my method was suitable and the experiment was successful because I had sufficient evidence to enable myself to come to a conclusion which agreed with my knowledge and prediction. I would have liked to share results with other people who were doing the same experiment as me to see if our results were similar, but nobody was doing the same experiment as me.

The only results which I did not think are reliable or accurate is the reaction of the 3. 8M concentration of hydrochloric acid with the calcium carbonate, and if I had more time I would investigate this further. I would find out why these results were anomalous because even though I did reapeats, I still got anomalous results and so I would like to find out why this happened. I would like to investigate the rate of reaction with more concentrations in order to see what happens after 3. 8M acid to see if it was still directly proportional or if the graph leveled off.

Other extra investigations I would do would include using different types of acid for instance nitric acid or sulphuric acid and see if they changed the reaction at all. I would also try and use different types of indigestion tablets, because the ones I used contained ginger which I have researched about to find that it is used for digestion, soothing aches and pains in muscles and improves circulation problems, so I would like to investigate if this has a different effect on the rate of the reaction or not.

I put an indigestion tablet into 1M of acid concentration and measured the temperature before and after the reaction for one minute to see if the reaction was exothermic or endothermic, but there was no change in temperature, so I would like to see what effect an exothermic or endothermic reaction would have on the experiment. Overall I think my results are reliable because the repeats are all very close to eachother, the biggest gap between my repeats is the 1. M concentration which had a time difference of 6 seconds, but the others which I had time to repeat are all around 3 seconds apart. If I had more time I would do much more repeats to make my results more reliable. The accuracy of my results are quite good because they are all very close to the line of best fit. I would like to do more experiments and repeats to make sure my line of best fit is accurate and in the correct place. Apart from these I think my investigation was successful.

DATE: September 16, 2012 TO: Patricia Bennett, Supervising Principal FROM: Connor Sims, Associate SUBJECT: Oil Drilling & Gas Extraction Industry in the US Analysis (21111) This report presents information regarding the industry, the primary operator of oil and gas field properties. The industry fuels its key buyers, the Natural Gas Distribution (22121) and the Petroleum Refining (32411) industries, with crude oil and natural gas. The industry continuously battles a shortage of available oil. In addition, many major oil fields have been in use for decades, slowly waning.

Currently, the industry grosses among the most profitable in the US despite these and similar obstacles. The benefits of investing here potentially outweigh concerning risks. Because of the esteemed value of the industry’s products, consistent demand for its products, and its positive near-future outlook, diversification into this industry may produce rewarding profitability in the short-term. High Product Value Crude Oil Prices The key economic driver for the Oil Drilling & Gas Extraction Industry, crude oil prices, determines much of its profitability according to supply and demand.

Price trends in West Texas Intermediate, a grade of crude oil used as a benchmark in oil pricing, display the growth of its value in the past 3 years and past decade. An average barrel of crude oil grew from $26. 18 in 2002 to $61. 95 in 2009, $79. 48 in 2010, and $94. 87 in 2011 (Airlines, 2012). JP Morgan analysts project average annual prices above $99 in upcoming years (Sethuraman, 2012). Such upward growth points to lucrative profits. Natural Gas Outlook Natural gas production accounts for 41. 6% of industry revenue in 2012. Prices n natural gas reached a 10-year low in April this year, but have erupted by more than 70% since (Hargreaves, 2012). Natural gas has seen an abundantly large output due to recent discoveries of natural gas in the Appalachian Basin; this large supply has kept prices relatively low recently, leaving opportunity for even higher profitability in future years. Consistent Demand Fueling US Industries The Oil Drilling & Gas Extraction Industry is the sole supply industry for its two demand industries, Petroleum Refining and Natural Gas Distribution (Hersch).

The US internally consumes 19,150,000 barrels of oil per day, doubling the world’s second largest consumer, China (Index, 2012). IBIS World describes the industry’s demand industries as “mature,” assuring the stable demand for our industry’s products (Hersch). Rising Exports, Foreign Buyers Current international relations appear conducive to this industry’s profitability. In 2011, for the first time since 1949, the US exported more refined oil than it imported (Winters, 2012); this evidences the success between the supply industry and its demand industries detailed above.

Additionally, oil exports to China will surge as it industrializes quickly. China’s exponentially growing demand leads to worldwide price increases (Hersch). Any increases, particularly increase this substantial, raise the WTI average price per barrel, increasing profitability. Positive Current Standing Favorable Market Concentration The four largest firms in the industry comprise of approximately 30. 0% of total revenue (Hersch, 2012). Market share concentration is low, allowing firms of any size to portion the industry’s $345. 9 billion revenue this year.

The competitive aspect of entering this industry would not be a difficult obstacle to overcome. Profit Margin The Oil Drilling and Gas Extraction Industry reels in a significantly larger margin in comparison to related industries. 46% of all industry revenue goes to profit, higher than the average for the entire mining sector, 39. 2% (Hersch, 2012). In 2008, the industry returned the 7th highest profit margin among US industries (Hargreaves). Profit margins have increased in the past 5 years as result of rising crude oil prices. Risks and Concerns Barriers of Entry

Most major oil and gas producers integrate services beyond drilling and extracting; many dualize as refining or distribution firms, circumventing demand industries en route to more direct profitability. New firms lacking this versatility may find an obstacle upon entry to the industry (Hersch, 2012). Additionally, firms in this industry must specialize in exploration and discovery for oil and gas resources. Firms may struggle finding initial success in this role due to the limited nature of resources. Long-term Resource Depletion ‘Peak oil’ refers to the prime of any field’s production, after which goes into terminal decline.

Most major US oil fields are beyond peak oil. The largest US oil field, Prudhoe Bay, has been depleting since 1979 (Prudhoe, 2012). The US Energy Information Administration indicates much production, particularly in the Alaskan North Slope, depends on world oil prices (Energy, 2012). Geophysicists and politicians debate over specifications regarding overall US peak oil, arguing the year in which US peak oil occurred. International Comparison In addition to the US peak oil situation, the US Oil Drilling and Gas Extraction Industry faces heavy foreign market competition.

In 2011, the US ranked 3rd in oil production, behind Saudi Arabia and Russia (Energy, 2012). Saudi Arabia’s OPEC governor expects Saudi output to rise steadily beyond 2030 with a 1. 5 million barrel per day spare production capacity then (Energy, 2012). Russia holds the world’s largest natural gas reserves, and its fuel exports have steadily increased since each year since 1999 (Energy, 2012). Conclusion Despite entry risks and threats of limited resources, evidence supports the likelihood of success for us to diversify into the industry under certain stipulations.

A new firm will implicitly face the challenge of exploring for land not already claimed by another firm. Additionally, alternative methods of energy will irrefutably have to replace oil drilling and gas extraction within an uncertain future; the remaining supply simply cannot match the demand forever. Two central obstacles hesitate immediate diversification: a barrier of entry and a negative long-term outlook. However, we must decide whether the benefits outweigh the concerns. World prices of oil and gas and China’s growing demand directly affect profitability.

Because evidence above shows substantial progress in both of these drivers with a very positive short-term outlook, diversification must be considered. If presence in the industry can be established quickly and will remain only until profitability falls, I recommend diversification.

References

1. Airlines For America (2012). Annual Crude Oil and Jet Fuel Prices. http://www. airlines. org/Pages/Annual-Crude-Oil-and- Jet-Fuel-Prices. aspx.. Retrieved September 16, 2012. Energy Information Administration (2012).
2. Project Alaska North Slope oil production at risk beyond 2025 if oil prices drop sharply. Today In Energy. http://www. ia. gov/todayinenergy/detail. cfm? id=7970
3. Retrieved September 16, 2012. Prudhoe Bay Fact Sheet (2012). British Petroleum. www. bp. com/assets/bp… us… /A03_prudhoe_bay_fact_sheet. pdf
4. Retrieved September 16, 2012. Hargreaves, Steve (2012). Natural gas prices surge 70%. CNN Money. http://money. cnn. com/2012/07/24/investing/natural-gas- – prices/index. htm.
5. Retrieved September 16, 2012. Hersch, Laura. (2012). IBIS World Industry Report 21111. Oil Drilling & Extraction In the US.
6. Retrieved September 16, 2012 from IBIS World Database. How the US Uses Oil (2012). Alternative Energy.
7. Retrieved September 16, 2012. ttp://alternativeenergy. procon. org/view. resource. php? resourceID=001797 Index Mundi (2012). http://www. indexmundi. com/g/r. aspx? c=us&v=91.
8. Retrieved September 16, 2012. Sethuraman, Nathan (2012). Poll: Increasing numbers see oil below $100 in 2013, 2014. Reuters. http://www. reuters. com/article/2012/06/27/us-oil-poll- idUSBRE85Q14720120627.
9. Retrieved September 16, 2012. Winter, Michael (2012). U. S. Exported more gasoline than imported last year. USA Today. http://content. usatoday. com/communities/ondeadline/post/2012/0 2/us-exported-more-gasoline-than-imported-last-year/1#. UFav7BhGhgI

An essential element of chemistry is finding reaction rates. This is because chemists need to know how long a reaction should take. In addition to needing to know the rate of a reaction at any point in time to monitor how the reaction is proceeding. Many factors effect reaction rates, two shown above include temperature and concentration. Concentration affects the rate of reactions because the more concentrated a solution the more likely collisions between particles will be.

This is simply because there are more particles present to collide with each other. When the temperature is higher, particles will have more energy. This means that more reactions will happen for two reasons, firstly more particles will come into contact with each other because they are moving around more and secondly because the reactions occur at higher speed making it more likely to succeed. A few other factors are the surface area and if a catalyst is present. The larger the surface area the more collisions will occur because there are more places for molecules to react with each other.

A catalyst affects the rate of reaction not by increasing the number of collisions, but by making more of the collisions that do occur successful. Ordinary household bleach is an aqueous solution of sodium hypochlorite, NaClO, this contains little more than 5% NaClO by mass. Bleaching is caused by the ion. Under normal circumstances this ion breaks down slowly giving off oxygen gas and the chloride ion, . In order to speed up this reaction a catalyst is needed. In this experiment the catalyst used was cobalt (II) nitrate solution.

When this is added to the bleach a black precipitate of cobalt (III) nitrate is formed which acts as a catalyst for the decomposition of The purpose of this experiment was to determine how concentration of reactants and temperature affect the rate of the reaction between bleach and 0. 01M cobalt (II) nitrate solution. In this experiment the volume of gas produced shows the rate of the reaction. Procedure Figure 1 Firstly, all safety protocols were ensured and applied (lab apron and safety goggles). The apparatus was set up with reference to figure 1 above.

Then, the eudiometer was filled with water and inverted into the trough, which was half filled with water. It was held in a vertical position with the burette clamp attached to the stand. The rubber tubing was joined to the top of the glass tube, which goes through the stopper on the flask. The other end of the tubing was then placed into the neck of the eudiometer. 15mL of bleach solution was measured into the 25mL-graduated cylinder and poured into the Erlenmeyer flask. As followed, 5mL of 0. 10M of cobalt (II) nitrate solution was measured and poured into the 10mL-graduated cylinder.

Once ready, the cobalt nitrate solution was poured into the flask containing the bleach solution, and the rubber stopper was immediately slotted in. It was then mixed and stirred as well as recorded (time). It was noted that a black precipitate of cobalt (III) oxide was forming, and from then on the flask was stirred gently and constantly. This was significant to dislodge bubbles of oxygen from the surface of the Co2O3 catalyst. Another thing that was important to note was that if the swirling was stopped or reduced, the rate decreases, so therefore the amount of swirling must be kept steady and uniform throughout the runs.

The total volume of oxygen that had been collected was recorded every 30 seconds until a volume of 50mL was obtained. Also, the actual elapsed time of when the 50mL mark was reached was recorded. Once the first run was successful, the following needed to be repeated the same way: the same amount of solutions must be measured into the same containers, and the procedure of applying them needed to be the same too (time recorded, measurements, temperature, etc. ). The only thing that was different in the next run was that the reactants had to be at a temperature of 10? C above room temperature before mixed.

This was accomplished by placing both the flask with bleach and the graduated cylinder with the cobalt (II) nitrate in a water bath for 10 minutes, and then adding the cobalt (II) nitrate to the flask, then back into the water bath. Hot water was used to increase the temperature, and cold water was used to adjust it. The next run was a similar idea to the previous one, but the reactants were brought down to a temperature 10? C below room temperature using ice. The steps to doing this are similar to the previous ones, but only this one required an addition of 20mL of water to the bleach solution before mixing.

The reason being is so that the overall concentrations are half of their original vales. The run that followed after was also identical, but instead of adding 20mL, 60mL was added. Now the overall concentrations after mixing were one quarter of their original values. The experiment was practically over, but there always had to be cleaning and instructed disposal of chemicals. The product(s) was/were instructed to be disposed in the designated container only for the waste solution. Finally, all the parties that participated in the experiment were obliged to wash their hands thoroughly with soap and water before leaving the laboratory.

Analysis and Results The rate of production of oxygen for each reaction was slightly different. The rate of reaction is determined by the equation; For the control where the reaction to place at room temperature and with bleach with a concentration of 0. 529M, the rate of production of oxygen was 36. 1 mL/minute. In next reaction which took place at a temperature 10? higher than that had a rate of 39. 5 mL/minute. Next was the reaction which took place at 10? below room temperature which resulted in a rate of 26. 8 mL/minute.

In the reaction that 20 mL of distilled water was added to the bleach solution and the temperature was kept constant, the reaction rate dropped to 16. 2 mL/minute. Finally the slowest reaction occurred when 60 mL of distilled water was added to the bleach causing a rate of 10. 8 mL/minute. The rate value changes as the temperature is changed. When the temperature increases by 10? , the rate of the reaction increases by a factor of 0. 12 (12%). This is again changed when the temperature is changed to 10? below room temperature. This results in a rate of production of oxygen, which is decreased by a factor of 0. 5 (25%). When the concentrations were changed so did the rate of reaction. When the concentration was changed to 0. 265M the rate of reaction dropped by a factor of 0. 5 (50%) below the control value. Furthermore when 60mL of water was added to the bleach dropping the concentration too 0. 132M the rate dropped by a factor of 0. 7 (70%). Bleach should never be mixed with any acid based cleaners because it results in the formation of toxic Cl- gas. If bleach is mixed with an acid based cleaner in a small room it will result in a toxic build up of chlorine gas, which can be fatal to anyone spending time in the room.

The equations for these reactions are shown below; Bleach is formed by the action of chlorine gas on sodium hydroxide, NaOH: The equation below represents the reaction of bleach with an acid based cleaner, which gives off chlorine gas Because of this reaction all acid based cleaners have warnings not to be mixed with bleach because it can result in injury or death. If bleach with 10% sodium hypochlorite was used for this experiment instead of bleach with 5. 25% sodium hypochlorite.

The shape of the rate curve for the graph would likely be twice as steep as the graph for the reaction involving bleach with a concentration of 5. 25% sodium hypochlorite. This is because the reaction will finish faster due a concentration that is higher by a factor of two. In this experiment there were possibilities for errors, the main one would be caused by measuring the volume of air at certain times. The reason for this being an error is that at 30 seconds more air will have been produced than is bing measure this is because some oxygen is in the Erlenmeyer flask but still rising to the point at which it is measured.

Also some oxygen is held back because of a kink in the rubber tubing. To improve this experiment I would use a better way of measuring the volume of oxygen produced, either by measuring the air pressure in a container attached to the flask where the reaction was taking place or by using a large tube with a piston inside that would slide along the inside of it showing how much oxygen is evolved during the reaction. Conclusion From the experiment that was carried out it can be concluded that both temperature and concentration effect reaction rates.

The lower the temperature the slower the reaction rate, therefore the higher the temperature the faster the reaction takes place. Likewise the lower the concentration of a solution the slower the reaction and the higher the concentration the faster the reaction is completed.

References

1. Measuring Reaction Rate Using Volume of Gas Produced. ” Experiment 11C. N. p. : SMG Lab, n. d. N. pag. Rpt. in Experiment 11C. N. p. : n. p. , n. d. 154-58. Print.
2. DiGiuseppe, et al. Reaction Rates. N. p. : Nelson, 2012. Print. Nelson Education.