The results of a chemical kinetics report from a simple chemical experiment

BASIC CHEMICAL PRACTICAL REPORT
KINETICS CHEMICALS

CHAPTER I

1.1. TITLE
Chemical Kinetics

1.2. DAY / DATE
 Thursday / May 18 2017

1.3. PRACTICAL OBJECTIVES
A. Measures changes in the concentrations of reagents by time
B. Observe the effect of concentration, temperature, and catalyst on reaction rate
C. Determining the rate laws of a reaction in an aqueous solution

1.4. PRAKTEK QUESTIONS
1. What is the concise definition of (a) rate law, (b) rate constant, (c) reaction order, (d) activation energy!
2. Is the unit of the reaction constant for (a). The order reaction is zero, (b). First-order reaction, (c). Second-order reaction!
3. Sulfur dioxide reduces HIO3 in acidic solution by reaction:
3iO2 (g) + 3H2O (l) + HIO3 (aq) à 3H2SO4 (aq) + HI (aq)
At the end of the reaction, if there is excess HIO3, this substance can be taken with a starch solution. The HI and HIO3 compounds immediately react to form I2 which is absorbed by the starch and gives rise to blue.
From the experiment obtained data:
[SO2] M
[HIO3] M
T (seconds)
14.6. 10-4
3.60. 10-3
25.8
7.31. 10-3
3.60. 10-3
52.8
14.6. 10-4
7.21. 10-3
12.6

Determine the order of reactions for each reactant and the whole order.
Answer.
1. (a) Rate law: right or left reaction can occur once the product is formed, this product can react again to produce all reactants.
(B) Rate stays: the provisions obtained from the experimental results and used to facilitate the process of finding the value of the rate
(C) Order of reaction: sum of all exponents of the concentration in the rate equation
(D) Activation energy: minimum kinetic energy required by the reactant particles to form the activated complex.
2. (a) zero-order reaction à k (mol s-1);
(B) first order reaction à k [A] [B] (s-1);
(C) second order reaction à k [A] 2 [B] 2 (mol-1 s-1)
3. Order SO2 Order of HIO3
    [] N = [] m =
             1.997n = 2 (0.5) m = 0.5
                                   N = 1 m = 1









CHAPTER II
2.1. Theoretical basis
Reaction rate measurement is an amazing chemical field. From kinetic studies, the mechanism of the reaction can be reduced. Information about the reaction of catalysts and inhibitors can only be obtained through kinetic assessment.
The rate of a chemical reaction can be affected by several factors:
· Reactant concentrations (and sometimes products)
· Temperature
· Catalyst
Rate measurements are usually performed under fixed experimental conditions, with one fixed factor whereas other factors are varied. When the influence of these factors on the rate has been determined, these factors are fixed and other factors are varied. Systematic assessment of rate dependence on reaction rate changes is continued until the kinetic behavior of the reaction is complete.
How to measure the rate of reaction, one important aspect of kinetic assessment is to design an easy technique to monitor the course of the reaction over time. Chemical analysis by means of volumetry or gravimetry is relatively slow, so this way is not in use unless the reaction is slow, or can be stopped by sudden cooling, or by the addition of reactants that stop the reaction.
Some commonly used ways are to use the nature of color and conductivity. The reaction rate involving the gas is determined by measuring the volume of time union gas. In this experiment you use a color change.

To a hypothetical reaction
2A + 3B à C + 5D
The law of its speed can be
Where k is the rate constant; N is the reaction order for A; And m is the reaction order for B. The overall reaction order is n + m. The order of the reaction can only be determined by experiment, since these numbers are not always the same as the reaction coefficient (stoichiometry).
In this experiment you will perform a reaction between sodium thiosulfate with hydrochloric acid.
S2O3-2 + 2H + à SO2 + H2
The rate of this reaction depends only on the concentration of S2O3-2, but not on the acid concentration. This is evidenced from the graph 1 / t against (S2O3-2) which gives a straight line. This graph implies that the order of reactions is one for thiosulfate. And because the acid concentration does not affect the rate. The rate of the org is zero.
You will also experiment with a two-order reaction for one of the reacting components. Second order can be inferred when a straight line is formed from graph 1 / t against [HCl] 2.
                                   (Basic Chemistry lab guide, 2011: 71-72)

Chemical reactions are the process of turning the reagents into reactions. The process is slow and there is fast. For example gasoline burns faster than kerosene. There are very fast reactions, such as burning dynamite that produces an explosion, and a very slow one is like a rusting process of iron. The discussion of the rate (rate) of the reaction is called kinetic chemistry. In this chemical kinetics put forward how to determine the rate of reaction and what factors influence it. (Syukri, 1999: 85)

Reaction kinetics is a branch of chemistry that discusses the rate of reactions and the factors that influence. The rate (velocity) of the reaction is expressed as a change in the concentration of reactant or reaction product to the unit of time. The reaction rate of a reaction can be expressed by the rate equation of the reaction, for the following reaction: A + B à AB The reaction rate equation is generally written as follows: R = k [A] m [B] nk as the reaction rate constant, m and n of the partial order respectively - the reactants. (Petrucci, 1987: 134)

The magnitude of the reaction rate is influenced by the following factors: a. Nature and size of reagents. The more reactive the reactant nature of the reaction rate will increase or the reaction progresses faster. The more surface area of ​​the reaction rate reactant substance will increase, this can be explained by the increasing surface area of ​​the reacting substance, the interaction of reactant substances is wider also. The surface of the reactant agent can be expanded by reducing the size of the reagent. So to increase the reaction rate, at the time of reactant in the form of powder is better when compared in the form of bongkahan.b. Concentration. From the general equation of the reaction rate, the rate of the reaction is proportional to the concentration of the reactant. If sodium thiosulfate is mixed with a strongly diluted acid, white precipitation will result. The reactions occur as follows: Na2S2O3 + 2H + à 2Na + + H2S2O3 (fast) H2S2O3 à H2SO3 + S (slow) Na2S2O3 + 2H + à 2Na + + H2S2O3 + This resonance consists of two consecutive reactions (continued connect). In such a reaction, a slow reaction determines the overall reaction rate. In this case the slowest reaction is the decomposition of H2S2O3.c. Reaction temperature. Almost all reactions become faster when the temperature is increased because the given heat will add the kinetic energy of the reactant particles. As a result the amount and energy of the collision grow larger. The effect of temperature change on the rate of reaction is quantitatively explained by Arrhenius law expressed by the following equation: k = or By R = ideal gas constant; A = typical constant for reaction (reactant factor) and Ea = activation energy in question. The catalyst is a substance added to a reaction to speed up the course of the reaction. The catalyst usually reacts temporarily and then re-forms as a free agent. A reaction using a catalyst is called a catalyst reaction and the process is called catalysis. The catalyst of a reaction is usually written over arrows (Petrucci, 1987: 81)

The order relates to rank in law of reaction rate. The constant reaction, not dependent on the concentration of the reactant is called the zero reaction order. The first-order reactions are more likely to exhibit a single concentration in the law of the rate, and the concentration is one rank. The most common formulation of the two-order reaction rate law is a single concentration of two or two concentrations of each of the first ranks. One method of determining reaction order requires measurement of the initial reaction rate of the experimental experiments. The second method requires precise mapping of the reactant concentration function over time. To get a straight line graph. (Hiskia, 1992: 213)

Based on the order of reactions, the reactions are differentiated into: 1. Zero-Order Reaction In zero-order reactions, the reaction rate is independent of the concentration of the reactants. The equation of the zero order reaction rate is expressed as follows: A - A0 = ko.tA = the concentration of the substance at time tA0 = Concentrations of the initial substance The example of this zero-order reaction is a heterogeneous reaction on the surface of the catalyst.2. First-Order Reaction In a first order reaction, the reaction rate is directly proportional to the concentration of the reactant. The equation of reaction rate of one is used as follows: If t = 0 à A = AoIn [A] = In [Ao] - k1t [A] = [Ao] The half-life (t1 / 2) is the time required for the reactant concentration to be only half . In a first order reaction, the half-life is expressed as follows. Second-order Reactions The equations for the second-order reaction are expressed as: Half-life for the second-order reaction is expressed as Reaction may be fast or slow. The factors that affect the fast and slow of a chemical reaction are: 1. Chemical properties of the reactants: in general ionic reactions are rapid, while reactions involving covalent bonds take place more slowly.2. The ability of interacting reactants: in the liquid or gas state of the reactant particles (molecules or ions) can collide easily with one another.3. Concentrations: the molecules must collide in order for a reaction in this context the reaction rate is proportional to the concentration of the reactant.4. Physical state: the molecules must mix with enough energy to react. Catalyst: The catalyst can be recovered without any chemical change. Temperature: molecule The current collides with enough energy to react. (Basic Chemistry II, 2005)

 CHAPTER III3.1. TOOLS AND MATERIALS · Tools1. Erlenmeyer2. Stopwatch3. Pipette drops4. Baja Sandpaper5. Pumpkin tip6. Tube reaction7. Glass Measure8. Watercolor · Materials1. Solution of Tiosulfate2. Air3. Hydrochloric Acid4. Na2S2O35. HCl6. Mg7 Band. Oxalic acid8. Sulfuric acid9. KMnO410. H2SO43.2. WORK PROCEDURES. Order reaction in sodium thiosulfate reaction with hydrochloric acid Reactivates reagents * Mixture Composition First mixed thiosulfate solution with water before Added hydrochloric acid Turned erlenmeyer to homogeneous mixture Recorded start time when acid is added to cloudy Formed again experiments, volumes of thiosulfate remain acidized volume Na2S2O3 (ml) [Na2S2O3 ] (M) H2O (ml) HCl (ml) 250.15-4200.1254150.09104100.0615450.03204 * ** Na2S2O3 (ml) H2O (ml) HCl (ml) HCl (M) 25-53.025231.825410.6B. Reaction order in reaction between magnesium with hydrochloric acid MgHCl 2MHasilCleaned with steel emulscrew Cut into 16 parts @ 2cmEnable 1 slice into erlenmeyerWatch the stopwatch when Mg dissolveMonsist Mg to keep movingDis timedPrice to erlenmeyer Diluted * in HCl (M) volume hoop (Ml) 0.61000.81001.01001 .21001.41001.61001.81002.0100 * C. Effect of temperature on reaction rateResults 3 drops KMnO4 0.1 Nair8 ml oxalic acid 0.1 N + 2 ml sulfuric acid 6N Added to 6 test tubes Incorporated into 3 glasses of glass 1 Glass: boiled Glass 2: heated Glass 3: not heated Inserted 2 tubes into each glass of trophies Added to each tube D. Effect of catalyst on KMnO4 reaction rate (3 drops) yield 6 ml as.oxalate + 2 ml H2SO46 ml as.oxalate + 4 ml H2SO46 ml as.oxalate + 6 ml H2SO4 Added to tubes 1 & 2 Inserted into tubes 3 & 4 Added to tubes 5 & 6 Dropped into each tube 3.3. DATA OF OBSERVATIONS. Reaction order in the reaction Sodium thiosulfate with hydrochloric acid Observations on the effect of Na-thiosulfate concentration Na2S2O3 (ml) Na2S2O3 (M) H2O (ml) HCl (ml) t (sec) 1 / t (sec-1) 250.15-41270,00787200.12541600,00625150.091041980, 00505100.061542860,0034950.032043190.000313Orde reaction is 1 for Na2S2O3 Observation on the effect of hydrochloride acid concentration Na2S2O3 (ml) H2O (ml) HCl (ml) [HCl] (M) t (sec) 1 / t (sec-1) 25-52.0780.012825231.8960. 010425410.61020.0098B. The order of reactions in the reaction of magnesium with hydrochloric acid Observations on the effect of HCl (M) HCl (ml) t (sec) 1 / t (sec-1) HCl 2 log (HCl) log (1 / t) 0.61004140, 0024150,36-0,221-2,6170.81003430,0028980,64-0,0960-2,5371.01002020,0045451,440-2,3421.21001300,007691,441.08-2,1141.4100880,011361,961.15-1,9451.6100550,018182,561.20- 1.7401.8100400,0253,241.25-1,6022.0100250,0441.30-1,398Order reaction is 1 for HclC. Effect of temperature on reaction rate Reaction time at various temperatures (sec) Deuteronomy Reaction temperature1000C500C250C12s13s33s22s13s33sRata - rata2s13s33sThe signs of reaction: a change of color to a clear whiteD. Effect of catalyst on reaction rateSlanganH2SO42ml1ml0ml12menit 33detik12min 4seconds23minutes 90detik22minute 41detik15menit24min 15secondsThe signs of reaction: a) 2ml → purple disappear (purple-purple-red liver-orange nodes) hot reaction) 1ml → purple disappear (purple-purple-red heart-orange Clear) heat reaction c) Purple color fixed, homogenous solution3.4. DISCUSSION. Order of reaction in reaction of sodium thiosulfate with hydrochloric acid In this experiment we observed the reaction order between the thiosulfate solution with hydrochloric acid. In the determination of the order of this reaction is closely related to the rate of reaction. In general, the rate of reaction is defined as the change in reagent concentration and reaction product per unit time. Chemically, the rate of the reaction is determined by determining the concentration of the substances at a given time, then the concentration data is used to calculate the reaction rate. In this experiment, we can determine the reaction order of each reaction reactor according to the reaction rate data. For the reaction: Na2S2O3 (aq) + 2HCl (aq) à2NaCl (aq) + SO2 (g) + S (s) In this experiment the thiosulfate solution is mixed with water, then hydrochloric acid is added. After homogeneous mixing, the experiments as listed in Table 10.1 and the results of the observations are as follows: a) 25 ml of Na2S2O3 with 0.15 M mixed with 4 ml of HCl. Within 127 seconds the solution has shown turbidity, due to the deposition of sulfur. In this first experiment the compound did not add water to the sodium thiosulfate. This included a faster mass appearing entirely in the solution when added with water or not with water.b) 20 ml of Na2S2O3 with 0.12 M mixed with 5 ml of H2O, then added with 4 HCl. Within 160 seconds the solution has shown turbidity, due to the deposition of sulfur at the bottom of the erlenmeyer tube.c) 1.5 ml of Na2S2O3 with 0.09 M mixed with 10ml H2O, then added with 4 ml of HCl. In 198 sThe solution has shown turbidity, due to the deposition of sulfur at the bottom of the erlenmeyer tube.d) 10 ml of Na2S2O3 0.06 M mixed with 15 ml of H2O, then added with 4 ml of HCl. Within 286 seconds the solution has shown turbidity, due to sulfur deposition at the bottom of the erlenmeyer tube.e) 5ml Na2S2O3 0.03 M mixed with 20 ml of H2O, then added with 4 ml of HCl. Within 319 seconds the solution has shown turbidity due to sulfur precipitation at the bottom of the test tube. In theory, for experiments with the mixed compositions listed in Table 10.2, we only look for the order of reactions alone and calculate the overall reaction order. Magnesium and hydrochloric acid. For reaction = Mg (s) + 2 HCl (aq) à MgCl2 (aq) + H2 (g) The treatments given in this reaction are: a) Comparison of various HCl concentrations (0.6M; 0.8M; 1.2M ; 1.4 M; 1.6M; 1.8M; and 2M). At the same volume of HCl and react the Mg band into the test tube. The observation result of bubble gas and magnesium dissolves in HCl until exhausted in a certain time.b) Magnesium takes several minutes to dissolve in HCl where speed depends on the concentration of HCl tersebut.c) The rate of reaction is influenced by the concentration of HCl vary. For example equation rate of reaction Formulated with r = k [Mg] [HCl], the reaction rate depends only on the phase of the [HCl] solution so that the reaction order to Mg is zero. The temperature effect on the reaction rate As previously known that the temperature rise accelerates the reaction, reaction. In the case of reaction rates: A2 + B à 2AB has the equation: r = k [A] m [B] nThe temperature change has k, since its value depends on the temperature and the type of reaction. If the temperature is increased, the amount and energy of collisions between reagent molecules increases. Effect of catalyst on reaction rate A slow reaction can be accelerated by giving / adding other substances without adding concentration or temperature. The substance is called a catalyst. The catalyst usually reacts temporarily and then re-forms as a free agent. Next reacts again with the reactant speeding up the reaction and free again. So on and so over and over again. A reaction using a catalyst is called a catalyst reaction and the process is called catalysis. The catalyst of a reaction is usually written over an arrow. Example: 2KCl (g) MnO2 2KCl (s) + 3O2 (g)


3.5. DISCUSION reactions in reaction Sodium Tiosulfate with hydrochloric acid In the observation experiment on the effect of Na-thiosulfate concentration, we used Na2S2O3 with different amount of volume, as well as H2O volume, different but HCl volume remained. This experiment succeeded, because we can prove the reaction order, where the reaction order x = 1. In addition, in subsequent experiments observation of the effect of concentration on hydrochloride acid, also obtained the order of reaction is 0.26. Thus we can calculate the overall order of the reaction is 1.26.a) Order of Reactions in the reaction of magnesium with hydrochloric acid In this experiment we observed the effect of hydrochloric acid concentration, just as our previous experiments used HCl with the same amount of volume, but the molarity was different . A total of 8 experiments. And the result of the experiment is that the higher the concentration of HCl used to react the Mg band, the faster the reaction rate.) The temperature effect on the reaction rate In this experiment, we have proven that the temperature affects the speed of reacting a substance, The sign of the reaction that occurs is the occurrence of changes in color. From the literature obtained, the temperature rise accelerates the reaction, and vice versa decreases the temperature slows down the reaction. In terms of reaction rates, for example, the reaction of A + B + C à result, has the equation V = k [A] m [B] n [c] 0The temperature change has k, since its value depends on the temperature and the type of reaction, is increased the sum and the collision energy Between reagent molecules increases. In addition, the reaction can take place in the event of a direct collision between reagent molecules. This collision must satisfy two conditions of effective position and sufficient energy. C) Effect of catalyst on reaction rate In this experiment, we tested the effect of the catalyst on the reaction rate of a solution and this experiment was successfully performed by the reaction signals is the occurrence of the color change.Reaction Which is slow to accelerate by giving other substances without adding concentration or temperature. This substance is called a catalyst. A reaction using a catalyst is called a catalyst reaction and the process is called catalysis. Based on the phase, the catalyst can be divided into two, namely homogeneous and heterogeneous cataracts.

3.6.QUESTIONS POST-PRACTICE1. Write the reaction equation on the experiment N C. Is H2SO4 in this experiment a catalyst? Explain! 2. Write the reaction equation in experiment D. Explain the working mechanism of H2SO4 as reaction catalyst! Answer: 1. The reaction equation: 5H2C2O4 (aq) + 2KMnO4 (aq) + 3H2SO4 (aq) à K2SO4 (aq) + 2MnSO4 (aq) + 10CO2 (g) + 8H2O (l) H2SO4 can act as a catalyst ie substances that can accelerate the reaction without Undergoes any changes.2. Equation of reaction: 5H2C2O4 (aq) + 2KMnO4 (aq) + 3H2SO4 (aq) à K2SO4 (aq) + 2MnSO4 (aq) + 10CO2 (g) + 8H2O (l) The working mechanism of H2SO4 as the catalyst in this reaction is: · H2C2O4 à CO2 + H2O · 2KMnO4 + 3H2SO4 à K2SO4 + 2MnSO4 Mechanism of catalyst work: P + R à PR (fast) PR + Q à PQ + R (fast) PQ: PereaksiP + Q + R à PQ + R (fast) R: CatalystH2SO4 increase Rate of reaction without any ingredients.

3.7. CONCLUSION From experiments conducted, the conclusions can be deduced: · The effect of time on the change of reagent concentration is the increasing of time, the reagent concentration is smaller. · The influence of temperature, concentration and catalyst on the reaction rate, the higher the temperature in the reaction, the faster the reaction rate . Similarly with concentration, the greater the concentration the faster the reaction rate, and the more the catalyst is introduced into a reaction the greater the concentration the faster the reaction rate. A reaction using a catalyst is called the catalyst reaction and the prosenya is called catalysis. Based on the phase, the catalyst can be divided into two: homogeneous and heterogeneous catalyst

3.8. REFERENCESE
three, dkk.2010.Unforcer Basic Chemistry Practicum.Jambi: Jambi Hiskia University, A and Tupamalu. 1992. Electrochemistry and Chemical Kinetics. Bandung: ITBPetrucci, Ralph H.1987. Basic Chemistry Principles and Applied Modern Volume 2. Jakarta: ErlanggaSyukri S, 1999. Basic Chemistry 2. Bandung: ITB