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International Baccalaureate: Chemistry
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Extraction is a major problem for aluminium and there is no ideal solution. However there are many not so perfect solutions that are currently being practiced today. In this essay I will look into some of the limitations and benefits of the current extraction techniques. I will also look at some of the economic and environmental effects the extraction of aluminium has. When aluminium is extracted from the other metals it must go through processes called electrolysis. Aluminium is not an easy metal to extract from other metals it is found in nature with, therefore it takes a lot of electricity to extract aluminium.
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Some examples of fuels that are released are carbon monoxide, nitrogen dioxide, nitric oxide. Different types of sulfur oxides and hydrocarbons are released from the combustion of these fossil fuels 2. When these gases are released into the environment, they cause smog and other signs of pollution. These are two major problems that are caused by burning fossil fuels. I will now look into a few alternative fuels that do not involve the burning of fossil fuels. I will look into each ones advantage and its limitations.
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Procedure- Variables- Independent Dependent Control The concentration of the acid Rate of reaction (volume of gas) Volume of HCL (1 mol l-1) Temp. (room temp) Mass of Calcium Carbonate(2.5g) Volume of Dilute HCL (50 cm3) Time (40 seconds) Agitation (none) Diagram- Materials- � 25 ml graduated cylinder � 100 ml graduated cylinder � Stopwatch � Stopper � Plastic tube � Side- arm test tube � Calcium carbonate with large(powder), medium(small chips)
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- ??HfO (Reactants) The standard enthalpy of formation is equivalent to the sum of many separate processes included in the Born-Haber cycle of synthesis reactions. Germain Henri Hess (1802 - 1850) is important primarily for his thermochemical studies. Hess' Law states that the heat evolved or absorbed in a chemical process is the same whether the process takes place in one or in several steps. This is also known as the law of constant heat summation. Hypothesis: The reaction being investigated is summarized in the chemical equation below: CuSO4(s)
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Descriptions of Different Chlorides. Silver Chloride Magnesium Chloride Lithium Chloride Potassium Chloride Calcium Chloride Sodium Chloride Ammonium Chloride Flame Colour Light Orange Orange Dark red Dark Orange Orange Light Orange Light Orange flames Substance description Solid, grey, chunky, neutral salt Crystal solution, clear, aqueous White crystal like powder, solid Chucky white crystals. Bigger than LiCl White solid blobs Solid white powder Solid white powder Quantitative Observations Wavelengths of Different Colours from Different Light Sources. Light Source Colour Daylight (�0.5 )
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Set up the beaker in the ring stand on top of the wire gauze 4. Measure about 20 ml of water with the graduated cylinder and place on top of the wire gauze in the ring stand. 5. Using the thermometer measure the temperature of the water and record it under control group. 6. Place the liquid fuel under the beaker and carefully light the candle 7. Using the stopwatch, start the time for one minute and blow out the candle after the minute has passed. 8. After you have blown out the candle, record the new temperature of the water.
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Heat again for a few minutes constantly stirring, then leave it to cool and reweigh 6. Repeat the procedure no 5 until there is no further loss in weight Drawing 1: The experiment Table 1: Variables DEPENDENT INDEPENDENT CONTROLLED The mass of salt The mass of water Intensity of heating Observations: Crystals of CuSO4*xH2O are blue due to water of hydration. When heated, the colour gradually changes from blue to gray-white. It's the effect of losing water molecules. The process can be easily observed on sides of crucible first. When the crystals are heated in temperature over 470 K, they lose all the water molecules.
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This would be done by using the oxidizing agent - potassium ferricyanide (K3Fe(CN)6). Control variables: * Concentration and volume of NaCl used * Concentration and volume of K3Fe(CN)6 used * Volume of phenolphthalein used Independent variables: * Iron (II) Sulphate FeSO4 crystal * NaOH solution * Iron nails - one bent and one straight Dependent variables: * Colour changes in the various reaction * Amount of rusting on the different nails Materials: Apparatus 1. 1 � measuring cylinder (50cm3) 2. 2 � dropper 3. 1 � beaker (50cm3) 4. 2 � test tubes 5.
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A metal can react with acid to form a salt and release hydrogen gas. A base can react with an acid to form a salt and water. An acid can react with a carbonate and form a salt, carbon dioxide and water. And finally two salts can react to form two other salts. AIM (PURPOSE) Our aim was to create 0,124 grams of calcium chloride. In order to prepare the salt, we used a base and an acid. The reason why we chose this method was first, we had the chemicals and second, using two salts was dangerous.
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The mass of the filter paper was measured and recorded. 3. A clean dry small beaker was put on the electronic balance and was tarred. 2.138 grams of sodium carbonate crystals were added and the exact mass was recorded. 4. A different beaker was then put on the electronic balance and was tarred. 1.040 grams of calcium chloride was added to and the exact mass was recorded. 5. Approximately 25 mL of deionized water was added separately to each of the beakers. Each beaker was stirred with different ends of a stir rod until the solids were dissolved. 6.
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Decreases from left to right (effective to nuclear charge) --> more protons and electrons b. Increases from top to bottom (valence electron shielding from the nucleus) - Ionization Energy a. Is the energy requires to remove an electron from an atom b. Increases from left to right c. Decreases from top to bottom because of valence (nuclear shielding) - Electron Affinity a. The ability for an atom to attract (or take) the electron away from another atom b. Increases from left to right c. Decreases from top to bottom because of nuclear shielding. Chemical Bonding - Isoelectronic: same electronic configuration - Ionic Bonds a.
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The solution of the Ca(OH)2 with the pink phenolphthalein indicator turned colorless when a certain amount of HCl (0.053mol/dm3) was added. Equation of reaction: 2HCl + Ca(OH)2 --> CaCl2 + 2H20 Moles = Concentration x Volume Moles of HCl = 22.72 x (0.053/1000) = 0.0012 Moles of Ca(OH)x = 0.0012/2 = 0.0006 (/2 Because of the 2:1 ratio between HCl and Ca(OH)2 To be able to make a comparison between the limewater concentrations of our obtained value and the literature value, we need to change it into g/100cm3 since the literature values are given in g/100cm3.
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In this case, we will use sample A. The Winkler method indicates that the relationship between the amount of oxygen and thiosulphate is a four to one ratio. Therefore, we will divide the number of mols of thiosulphate by four to get the number of mols of oxygen. Then we convert this to milligrams per dm� first by multiplying by 20 to arrive with the number of mols present in 1dm� (because the pond water sample used was 50cm�). Then we multiply that number by 32 (molecular mass of oxygen)
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22.10 7.4 26 22.20 8.3 27 22.20 8.4 28 22.30 8.4 29 22.40 8.6 30 22.50 9.0 31 22.70 9.5 32 22.80 9.6 33 22.90 10.0 34 23.00 10.1 35 23.20 10.2 36 23.40 10.6 37 23.60 10.6 38 23.70 10.7 39 23.80 10.8 40 24.00 10.9 41 24.20 10.9 42 24.40 10.9 43 24.60 11.1 44 24.80 11.0 45 25.00 11.0 46 25.20 11.1 47 25.40 11.2 48 25.60 11.1 49 26.00 11.2 50 27.00 11.3 51 28.00 11.4 52 30.00 11.4 53 32.00 11.5 54 35.00 11.6 55 40.00 11.7 56 45.00 11.7 57 50.00 11.7 Weak acid- Strong base titration: Both NaOH (Strong base)
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Before adding water Magnesium oxide had a white appearance resembling bird droppings. This changed it to grey color like a cigarette ash. Percentage Error = (moles of Oxygen- moles of Mg) x100% Moles of oxygen = (6.3 X10- 6.7 X10) x100% 6.3 X10 = 0.6% Conclusion and evaluation: The empirical formula obtained experimentally from the lab defers from the theoretical empirical formula, proving that the results of the experiment do not support the theoretical assumptions about the formula for magnesium oxide.
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These values are presented in Table 3. For the experiments that were taken with 2 trials, an average time was used. The time was inversed in order to graph ln k = -(Ea/R)(1/T) + ln A Average time sample calculation: (676.4 s +710.9 s) / 2 = 693.65 s Sample calculation of k (at 2.5 �C) k = Initial Rate of Reaction [KMnO4][H2C2O4] k = 9.50977129 x 10-6 mol dm-3 s-1 (0.02 mol dm-3)(0.5 mol dm-3) k = 9.50977129 x 10-4 mol-1 dm-3 s-1 Table 2 - Determining Average time and Initial Rate of Reaction Temperature (�K � 0.2� K)
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I intended to approach this problem by using my knowledge of chemicals stated above to come up with a simple, practical way of concocting five grams of Barium Sulfate. Hypothesis If Barium Nitrate and Sulfuric Acid are mixed together appropriately, then we will be able to make five grams of Barium Sulfate, along with the other product. This hypothesis is clearly justifiable. If you mix 5.6 grams of Barium Nitrate with 2.1 grams of Sulfuric Acid, through the reaction types we can see that this is double replacement and should form HNO3 + BaSO4, and through some dilution calculations we can see that theoretically five grams of Barium Sulfate should be made.
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Dependent o The changes in temperature for the products, or the ?enthalpy (?H) were the dependent variables for this lab. Control o The beakers, graduated cylinder, stopwatch gloves, goggles, and the times recorded were the controls for this lab. Part 2 Independent * The amount of reactants needed (7.1 grams of diluted HCl and 11.7 grams of diluted KOH) and the amount of water needed to dilute them into their solutions were the independent variables. Dependent * The changes in temperatures needed to find the actual enthalpies to compare to our theoretical values.
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less expensive zinc used as a filter in the US pennies made after 1983 versus the percentage of the almost pure copper used previous to 1983. Procedure: 1. Obtain pennies that are minted after 1983 and cut them into half-sections using metal shears. Note the silver-colored zinc core and the outer copper layer 2. Record the mass of two of the cut penny sections to the nearest 0.001 g. Place them into a labeled 100 ml beaker. 3. Repeat step 2 with the other remaining half-section. 4. To each beaker, add about 20 mL of concentrated (12 M) hydrochloric acid.
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The kernels in the beaker were heated using a Bunsen burner. (See Figure 1 - Setup Diagram) While the kernels popped, the beaker was shaken in order to prevent the popped popcorn from burning and to shake the unpopped kernels to the bottom. Once all the kernels had popped, the beaker was removed from the Bunsen burner and set to cool. After it had cooled, the beaker with foil, oil, and kernels were massed again. The same steps were repeated for the rest of the samples and qualitative observations were made throughout the experiment.
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Catalyst - no catalyst is added to lower the activation energy of the reaction. The reaction can therefore not be accelerated because it would otherwise affect the rate of reaction. Surface area is kept the same by keeping the same size of beaker over the marked cross when mixing both chemicals together. Reasons why controlled variables would have an effect on the rate of reaction if they weren't controlled: Condition Effect on rate Explanation Temperature increasing the temperature increases the rate of a reaction Two reasons: 1. There are more particles with sufficient energy to react (most important)
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What was supposed to happen was that once the two liquids were mixed, after stirring for no more than 5 minutes, the solution was to turn clear, transparent, and colorless. However, in this trial, there was still quite a strong purple color even after of 10 minutes of mixing. By this time, I realized that the liquids had cooled down, and the reaction was going to take a lot longer to undergo, if at all. Thus, the next time I ran the experiment at this temperature, I left the beaker in which I was mixing the solutions on top of the hot water bath, thereby keeping it heated, and the reaction did eventually take place.
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Air bags are designed to complement conventional restraints such as seatbelts and seatbelt pre-tentioners, not replace them. In this experiment, a simple airbag will be constructed which will inflate on an impact. The bag will be made of thin plastic. A chemical reaction will be used to inflate the airbag. HYPOTHESIS: Reaction of an acid with a carbonate salt should liberate enough gas, in a closed bag, to make an airbag. VARIABLES: * Independent variables - Quantity of reactants used. * Dependent variables - The amount of gas released to fill the pouch. * Controlled variables - Temperature and pressure.
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The mass of cold water. m(water)=?(water)�V(water) =1g/mL�50mL =50g Absolute uncertainty=�1mL (Assume that the absolute uncertainty of ?(water) is negligible) % uncertainty=1/50 � 100%=2% b) The change of the temperature ?T is the temperature change after the warm water was added into the cold water. ?T=38.2?-20.2? =18? Absolute uncertainty= �0.5? + �0.5?=�1? % uncertainty=1?/18? � 100%=6% c) The heat absorbed by the cold water Q(cold water)=m�s�?T =50g�4.18J/g?�18? =3.8�103J =3.8kJ % uncertainty=2% + 6%=8% (Assume that the absolute uncertainty of the specific heat of water is negligible)
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EQUIPMENT & MATERIALS * * Safety goggles/Lab coat * Four Fruit juices (100 mL) * Vitamin C standard solution (1.0 g/mL) * Starch solution (1%) * Iodine solution * Hydrochloric acid (HCl) 1M * Distilled water * 50-mL burette * Burette * Retort Stand and Burette Clamp * 10ml Graduated pipette and pumpette * 25-mL Graduated cylinder * Two 125-mL conical flasks * Two 100-mL beaker PROCEDURE 1. A 10-mL graduated pipette was used to obtain 10.0 mL of the vitamin C standard solution, and placed in a flask.
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