Calculate; a. The average rate of the reaction. b. The rate of reaction at 2 ... The following balanced chemical equations represent the chemical reactions of the ...
College of Natural and Computational Science
Stem manual for grade 11
Prepared by:
Belete kefarge (MSc. In Materials Science) Angaw kelemework (MSc. In Inorganic Chemistry) July, 2014 i
Table of Contents Unit 1: Chemical bonding ............................................................................................................... 1 Experiment 1: Electrical Conductivity of Ionic Compounds ...................................................... 1 Unit 2: Rates of Chemical Reaction ................................................................................................ 3 Experiment 2: Measurement of a Reaction Rate ........................................................................ 3 Experiment 2.1: Effect of Concentration on Reaction Rate........................................................ 7 Experiment 2.2: Effect of Nature of Reactants on Reaction Rate .............................................. 8 Experiment 2.3: Effect of Temperature on Reaction Rate ........................................................ 10 Unit 3: Chemical Equilibrium ....................................................................................................... 12 Experiment 3.0: Determination of Equilibrium Constant for an Organic Acid ........................ 12 Experiment 3.1: Effect of Change in Temperature on the Equilibrium position .................... 16 Experiment 3.2: Effect of Change in Concentration on Equilibrium Constant ........................ 17
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Unit 1: Chemical bonding Experiment 1: Electrical Conductivity of Ionic Compounds Objective: To test the electrical conductivity of aqueous solutions of ionic compounds Theory: Chemical Bonding The foundations of modern chemical bonding theory were laid in 1916–1920 by G.N. Lewis and I. Langmuir who suggested that ionic species were formed by electron transfer, while electron sharing was important in covalent molecules. In some cases, it was suggested that the shared electrons in a bond were provided by one of the atoms but that once the bond (sometimes called a coordinate bond) is formed, it is indistinguishable from a ‘normal’ covalent bond. In a covalent species, electrons are shared between atoms. In an ionic species, one or more electrons are transferred between atoms to form ions When an atom either loses or gains electrons, it becomes an ion. An ion is an electrically charged particle. There are two types of ions. They are the positive ions called cations and the negative ions called anions. A positive ion is formed when an atom (a metal) loses one or more electrons. A negative ion is formed when an atom (a non-metal) gains one or more electrons. When a cation and an anion are brought close to one another, a force of attraction is set up between them. This force of attraction between oppositely charged ions is called an ionic bond. Compounds which are formed by an ionic bond are called ionic compounds. Ionic compounds show some physical properties that are common to all. One such a property is their ability to conduct electricity. The following experiment illustrates the electrical conductivity of ionic compounds. Ionic bond :A type of chemical bonding in which one or more electrons are transferred completely from one atom to another, thus converting the neutral atoms into electrically charged ions; these ions are approximately spherical and attract one another because of their opposite charge. Also known as electrovalent bond.
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Apparatus 9 -Volt battery with a bulb holder 6 -Watt bulb with a bulb holder Conducting wires Two carbon rods, Beaker Chemicals Water(H2O) Lead iodide (PbI2(aq)) or Sodium chloride ( NaCl (aq)) Procedure 1. Dissolve some amount of PbI2 or NaCl in 50 mL of water in a beaker. 2. Connect the circuit as shown in Figure 1. 3. Test the conductivity of the solution. 4. Repeat the procedure for pure water and aqueous solution of sugar.
Figure 1 Electrochemical cell showing the conductivity of an aqueous solution. Observations and analysis: The bulb lightens in aqueous solution of PbI2 or NaCl . It shows that the ionic solution conducts electricity. The charge carriers are Na+ and I- ions (or Na+ and Cl-). The bulb does not lighten in pure water and in aqueous sugar (covalent compound) solution. Conclusion: Solutions of ionic compounds conduct electricity while covalent compounds do not conduct electricity 2
Unit 2: Rates of Chemical Reaction Experiment 2: Measurement of a Reaction Rate
Objective: To measure the rate of a reaction between CaCO3 and dilute HCl Theory: Rates of Chemical Reaction Introduction Chemical kinetics is the branch of physical chemistry concerned with the mechanisms and rates of chemical reactions also known as reaction kinetics. It is concerned with the speeds, or rates, at which reactions occur. Chemical kinetics refers to the rate of a reaction. The rate of a chemical reaction is a measure of the decrease in concentration of a reactant or the increase in concentration of a product per unit time. The rate of a chemical reaction is obtained by determining the concentration of reactants or products at certain time intervals during the reaction, mathematically: Rate of reaction =
chargein concentration of a substance chenge in time
∆C
= ∆t
From the above mathematical expression, it follows that the rate of a reaction is inversely proportional to the time taken by the reaction and directly proportional to the concentration of the substance. Accordingly, the rate of a reaction decreases with increasing reaction time. Figure 2 below illustrates the change of the rate of a chemical reaction with time.
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Figure 2 The change in concentration of product with time. The methods for determining the concentration of reactants or products depend on the type of the reactions. Some of the methods of determining the concentration of reactant and products are: 1. Changes in color 2. Changes in pressure particularly for gases 3. Changes in volume particularly for gases 4. Changes in weight 5. Amount of precipitate formed The factors that affect the rate of a chemical reaction are: 1. The nature of reactants: More reactive elements react faster. 2. Temperature: Temperature usually has a major effect on the rate of a reaction. Molecules at higher temperatures have more thermal energy. Generally, an increase in the temperature of a reaction mixture increases the rate of the chemical reaction. This is because as the temperature of the reaction mixture raises, the average kinetic energy of the reacting particles increases. So they collide more frequently and with greater energy. The effect of temperature is experienced in our daily life. We keep our food in a refrigerator so as to slow down its rate of decomposition reaction. 4
3. Concentration of reactants: The higher the concentration of the reactants, the higher the rate of the chemical reaction. 4. Surface area of reactants: the higher the surface area of the reactants, the faster is the rate of the reaction. 5. Catalysts: a catalyst speeds up the rate of a reaction by providing an alternative reaction path with lower activation energy. In the following experiments the measurement of the rate of a chemical reaction is illustrated. Furthermore the effect of concentration, nature of reactants and temperature on reaction rate is investigated. Apparatus Analytical Balance 100 mL conical flask Stop watch Cotton wool. Chemicals 1M dilute hydrochloric acid( HCl )solution Calcium carbonate (CaCO3 )
Procedure: 1. Setup the apparatus as shown in Figure 3. 2. Add 20 g CaCO3 in a clean 100 mL conical flask. 3. Add 40 mL of 1 M hydrochloric acid to the conical flask. 4. Plug the cotton wool in position immediately. 5. Read the mass of the flask and its contents and start the stop watch. 6. Record the mass at one minute intervals for ten minutes.
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Figure 3 Laboratory setup for the measurement of reaction rate.
Observation and analysis: 1. Use the following Table to record your observations.
Time (min.) 0
1
2
3
4
5
6
7
8
9
10
Mass (g) Decrease in mass (g)
Plot a graph with time (minutes) on the horizontal axis and change in mass on the vertical axis. Draw a smooth curve through as many points as possible. Note that in this experiment change in mass is proportional to change in concentration.
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2. The mass decreases during the reaction, because one of the reaction products, CO2 gas, is leaving the system as shown in the following reaction. 2HCl + CaCO3 → CaCl2 + H2O + CO2↑
3. You will observe that the graph is steeply (hence high change in mass implying fast reaction rate) in the beginning. Slowly, the steepness decreases (hence small change in mass implying slow reaction rate) and becomes horizontal at the end of the reaction. This shows that the rate of the chemical reaction decreases with increasing reaction time. This is because the concentration of the reactants goes on decreasing with increasing reaction time. 4. Calculate; a. The average rate of the reaction. b. The rate of reaction at 2 minutes.
5. Can we measure the rate of the reaction by measuring the amount/volume of CO2 evolved
Experiment 2.1: Effect of Concentration on Reaction Rate
Objective: To study the effect of concentration of hydrochloric acid on the rate of its reaction with CaCO3 Apparatus digital balance 100 mL conical flask stop watch cotton wool Chemicals HCl solutions marble chips (CaCO3)
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Procedure Repeat the procedure in experiment 2.0 but use 0.5 M HCl and 2 M HCl solutions. Observations and analysis 1. Compare the rate of the reactions when using 0.5 M HCl and 2 M HCl solutions. Which reaction is more vigorous? Explain. 2. Which of the two reactions will produce more carbon dioxide? Explain your answer. 3. Compare the rate curves and average rates when 0.5 M, 1 M and 2 M HCl solutions react with CaCO3.
Experiment 2.2: Effect of Nature of Reactants on Reaction Rate
Objective: To study the effect of nature of reactants for the reaction of copper and magnesium metals with hydrochloric acid Apparatus Balance Test tubes Test tube stand Eye goggles Gloves Chemicals Copper Magnesium 2 M HCl Procedure 1. Measure equal masses of copper and magnesium metals using a balance. 2. Record the weighed mass of these metals and add each to separate test tubes. 3. Assemble the test tubes as shown in Figure 4. 4. Add 5 mL of 1 M HCl to each of the test tubes. Be sure to wear your safety goggles and gloves. 5. Observe relative rates of reaction in both test tubes and record your observation.
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Figure 4 Reaction of magnesium with HCl. Observation and analysis: 1. You have observed that magnesium reacts faster than copper. It is because magnesium is more reactive than copper. 2. The following balanced chemical equations represent the chemical reactions of the two metals with HCl. Mg + HCl → MgCl2 + H2↑ Cu + HCl → CuCl2 + H2↑
3. The experiment shows different reactants react with a given compounds at different rates depending on the reactivates of the reactants. More reactive reactants react faster while less reactive reactants react slowly. 4. If similar experiment is done using zinc and aluminum metals, which reaction do you expect to be faster? Why?
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Experiment 2.3: Effect of Temperature on Reaction Rate
Objective: To study the effect of temperature on the rate of the chemical reaction between sodium thiosulphate (Na2S2O3) and hydrochloric acid (HCl) Apparatus 100 Ml beakers Test tubes Thermometer white paper pencil water bath Chemicals 0.5 M HCl 0.1 M Na2S2O3 pieces of ice Procedure 1. Take 25 mL of 0.1 M Na2S2O3 solution in a test tube and 25 mL of 0.5 M HCl solution in another test tube. 2. Prepare 3 such sets and maintain them at different temperatures. Set a at 0oC (by keeping them in an ice bath as shown in Figure 5. Set b at room temperature. Set c at 40oC (by heating the two solutions in a water bath). 3. Put a cross sign on a white cardboard and place a clean dry 100 mL beaker above it. 4. Now pour the contents of set a in the beaker and start the stopwatch immediately. 5. Carefully stir the mixture with a glass rod and record the time taken for the cross sign to disappear and Repeat steps 3, 4 and 5 with set b and set c, respectively. 6. Tabulate your results as temperature (in OC) versus time (in minute).
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Figure 5 Laboratory setup for the study of the effect of temperature on reaction of Na2S2O3 & HCl. Observation and analysis 1. Sodium thiosulphate reacts with hydrochloric acid producing sulphur, sulphur dioxide, sodium chloride and water. Na2S2O3 + HCl → S + SO2 + NaCl + H2O The sulphur produced forms small particles and causes the solution to be cloudy and turn to a yellow color. This causes the cross sign to fade and eventually disappear. The time needed for the cross sign to disappear is used as a measure of the reaction rate. Faster reactions need less time for the cross sign to disappear. 2. Plot the graph of time (minute) taken for the cross sign to disappear on the horizontal axis against temperature on the vertical axis. A linear relationship is expected. Conclusion o A reaction rate increases with increasing temperature.
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Unit 3: Chemical Equilibrium Experiment 3.0: Determination of Equilibrium Constant for an Organic Acid
Objective: To determine the equilibrium constant for the esterification of acetic acid. Theory Esterification is one of the characteristic reactions carboxylic acids or their functional derivatives (acid chlorides and acid anhydrides) undergo with alcohols to produce esters. In the presence of catalytic amount of an acid, the direct esterification reaction involving a carboxylic acid and an alcohol, also known as Fischer esterification, proceeds reversibly according to the following general equation.
The equilibrium constant, Kc, for this reaction is given by the following expression: Kc = [Acid] [Alcohol]/[Ester] ,where the terms in square brackets represent the equilibrium concentration of unreacted acid, alcohol and ester product. The magnitude of the equilibrium constant Kc indicates the extent to which the reaction has proceeded at a given temperature. The yield of the ester product in this equilibrium reaction can be raised by application of Le Chatelier’s principle: judicious decision to use an excess of one of the reactants (alcohol or carboxylic acid), or continuously remove one of the products (ester or water) from the reaction mixture. Typical acid-catalyzed esterification reactions between acetic acid (or ethanoic acid) and alcohols butan-1-ol and propan-1-ol is depicted in the following equation.
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Acid anhydrides and acid chlorides also react with alcohols at a rather faster rate than the reactions involving the corresponding carboxylic acids to produce esters according to the following equations.
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Apparatus Round bottom flask Pipette Conical flask Burette Burette stand Water condenser with inlet and outlet pipes Boiling chips Clamp stand Bunsen burner Water bath. Dropper Chemicals Acetic Acid Propan-1-ol Phenolphthalein Sodium Hydroxide Solution Procedure: 1. Into a 100 mL round bottom flask attached to a stand with the help of a clamp, add 10 mL of acetic acid and 10 mL of propan-1-ol 2. Add cautiously about five drops of concentrated sulfuric acid into the flask. Swirl gently the flask with your hand to ensure mixing of the contents in the flask. 3.
Introduce a couple of boiling chips into the flask.
4.
Attach a water condenser to the flask as shown in Figure 1 below.
5. Open the water tap to allow water flow through the condenser and reflux for 1 hour. 6. Cool the flask and its contents in an ice bath. 7. Prepare ice cold water by placing 25 mL of water in a conical flask and letting the conical flask stand on an ice bath for several minutes. 8. Pipette out 1.0 mL of the reaction mixture and add it to the conical flask containing the ice cold water. 9. Add 2 drops of phenolphthalein indicator to the solution in the conical flask.
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10. Titrate the contents in the conical flask against a 0.1 M solution of sodium hydroxide. 11. Repeat the last three operations pertaining to titration three times. Note down all your observations. Exercise: 1. Why is sulfuric acid added to the reaction flask at the beginning of the experiment? 2. Calculate the weights of acetic acid and propan-1-ol used in this experiment. 3. Calculate (a) the concentrations reactants and products at equilibrium, (b) the equilibrium constant for the esterification of acetic acid.
Figure 6. A reflux setup
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Experiment 3.1: Effect of Change in Temperature on the Equilibrium position
Objective: To determine the effect of change in temperature on the position of equilibrium of iodine and starch. Theory The conversion of a given set of reactants into product(s) may occur quantitatively. That is, the reaction may go essentially to completion. Such reactions are said to be irreversible. On the other hand, many reactions are known that proceed reversibly. That is, the reaction attains a state of dynamic equilibrium wherein the forward and reverse reactions occur at equal rates. This balance in rates between the forward and reverse reactions can be disturbed temporarily by the following factors: a. Change in concentration of a reactant/product b.
Change in pressure and volume of the system, and
c. Change in temperature. A catalyst, however, has no discernible effect on the position of equilibrium. The effect of such changes imposed on a chemical reaction that has attained a dynamic state of equilibrium can be predicted by Le Chatelier’s principle: If a chemical system at equilibrium experiences a change in concentration, temperature, or volume, then the position of the equilibrium shifts so as to counteract the imposed change thereby establishing a new equilibrium position. For instance, for the following hypothetical reaction at equilibrium, if the concentration of reactant M is increased, the system will respond by consuming more of A through reaction with reactant N thereby producing more of products X and Y. That is, the system responds to the stress by pushing the equilibrium position in favor of the products. M + 2N
X+Y
The value of Kc for a given reaction will change only with a change in temperature. The effect of a change of temperature on a reaction will depend on whether the reaction is exothermic or endothermic. As per Le Chatelier's principle, increasing the temperature of an endothermic reaction will drive the reaction in the forward direction while the same effect would favor the reverse process for an exothermic reaction. On the other hand, raising the temperature of an exothermic reaction would assist the reverse reaction while lowering the temperature of same reaction will assist the forward reaction.
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In this experiment, you will observe the effect of temperature on the equilibrium involving starchiodine complex. A very low concentration of iodine solution (iodine dissolved in an aqueous solution of potassium iodide) reacts with starch producing an intensely colored starch/iodine complex. Apparatus Test tubes water bath, Stands Bunsen burner Thermometer. Chemicals Iodine Starch Procedure: 1. Into a clean test tube containing some starch solution, add a few drops of tincture of iodine. Note down what happens. 2. Place the test tube on a water bath and heat the solution to about 800C and record your observations. 3. Cool the test tube in an ice box. Record your observations. Activities Describe the effect of temperature on the starch-iodine complex forming reaction.
Experiment 3.2: Effect of Change in Concentration on Equilibrium Constant
Objective: To study the effect of change in concentration on the position of the equilibrium for the reaction between Fe(NO3)3 and KSCN Apparatus Test tubes Test tube rack Beakers (100 mL capacity)
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Chemicals 0.05 M Fe(NO3)3 solution 0.01 M KSCN solution 0.01 M HNO3 solution. Procedure 1. Take 5 test tubes that are clean and dry. Label the test tubes 1, 2, 3, 4, and 5 and place them in a test tube rack. 2. Take 50 mL of 0.05 M solution of Fe(NO3)3 in a 100 mL beaker. In separate beakers, take 10 mL of 0.01 M KSCN solution and 20 mL of 0.1 M HNO3 solution. Mix the solutions as per the specifications given in the table below to prepare five different solutions.
Test tube No
Volume in mL of 0.05 M Fe(NO3)3
0.1 M HNO3
0.01 M KSCN
1
1.0
4.0
1.0
2
2.0
3.0
1.0
3
3.0
2.0
1.0
4
4.0
1.0
1.0
5
5.0
0.0
1.0
Note the color of the solution in each test tube and arrange the test tubes in increasing order of color intensity. Activities You have studied the following reaction at a chemical equilibrium. Fe+3(aq)
+
SCN- (aq)
Fe(SCN)+2(aq) deep red
1. Which direction is exothermic and which is endothermic? How do you explain your results? 2. Compare the color intensity with the concentration of Fe+3 in the test tube. (Hint: As shown in above equation, Fe+3 forms deep red colored complex with SCN-.)
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3. By taking different volumes of Fe(NO3)3 in the test tubes the concentration of Fe+3 is varied in the solution. Calculate the concentration of Fe+3 ions in each test tube and correlate with the color intensity. 4. In the test tube take 1.0 mL of Fe(NO3)3 solution and add 4.0 mL of HNO3solution followed by 1.0 mL KSCN solution. Mix well and note the color. Add 1.0 mL of Fe(NO3)3 solution and again note the color of the solution.
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