The Effect of Concentration on the Reaction Between Magnesium and Hydrochloric Acid

The Effect of Concentration on the Reaction Between Magnesium and Hydrochloric Acid
IntroductionChemical kinetics is the study of how fast (rate) a reaction takesplace. We define reaction rate as the change in concentration of asubstance divided by the time taken for that change to take place(Johnson et al., 1996) its units are mol dm-3/s. Without theunderstanding of kinetics scientists wouldn?t understand basicchemical life processes, and how complications occur. For example,high levels of oxygen need to be maintained within cells, oxygen is inconstant demand; its interaction with molecules results in chemicalreactions necessary for life. If concentrations of oxygen decrease,rate of reactions decreases, resulting in impairment of cell function,and death (Seeley et al., 2000).Many factors influence the rate reactions proceed, how easilysubstances react with one another for example, all reactants differ inabilities e.g. iron corrodes faster than stainless steel (Parsons etal., 2002).
Other factors include concentration and surface area of reactants, presence of catalyst and temperature. An example of the effect of temperature can be seen in a person with a fever, increases in temperature cause intracellular reactions to accelerate, resulting in increased activity within organ systems e.g. increased heart rate. The aim of this experiment is to determine if the reaction rate between Magnesium (Mg) and Hydrochloric acid (HCl) is affected by varying concentrations of HCl. Balanced equation: Mg(s) + 2HCl(aq) → MgCl2(aq) + h2 Within limits, the graeter the concentration of a reaction, the graeter the rate at which a chemical reaction proceeds (Seeley et al., 2000). This occurs as the molecules are more likely to come into contact with each other, also known as the collision theory of reactivity. Reactions occur due to molecular collisions, during collisions, bonds are broken and new molecules are re-formed (Ritchie et al., 2003). Not all collisions result in a successful reaction however, molecules may just recombine with their ?previous partners?. Collisions require energy; strong electrostatic forces need to be overcome to separate molecules. Insufficient energy results in no reaction, molecules just bounce off one another. Therefore, it can be assumed that if the concentration of HCl/Mg is increased, more collisions will occur, resulting in faster reaction rates. If more particles of HCl/Mg are present, the probability of a collision occurring increases, the number of collisions will exceed the activation energy resulting in a rise. Method Experiment 1, Magnesium as the subject. 6 lengths of Mg ribbon were cut (2cm-7cm), weighed (to calculate moles), and cleaned with emery paper. Six tests were carried out whereby each piece was placed in an Erlenmeyer flask, excess (50cm3) HCl (2M) was added followed promptly with a bung. The amount of H 2 produced was collected and recorded using a gas syringe. Readings were taken at 5 second intervals over 30 seconds. The amount of Mg to use in experiment 2 was established (2.1×10-3). This amount was decided as a constant, noticeable increase in the amount of H2 produced was shown. The amount was neither too low nor high and was easily recorded. Experiment 2, Hydrochloric acid as the subject. A 2 Molar (M) solution of HCl was diluted to obtain four other concentrations (1.6, 1.2, 0.8, 0.4 M). The same method was carried out as in experiment 1, except the amount of Mg was kept constant throughout. 50cm3 of each HCl dilution was used. All experiment was carried out at room temperature and pressure. Results Experiment 1. [Mg] + HCl (2M). Table 1. Volume of H2 produced for n* moles of Mg + HCl (2M) over 30 seconds [Mg](moles) â?? 8.33×10-4 1.25×10-3 1.66×10-3 2.1×10-3 2.5×10-3 2.92×10-3 Time(sec)â?? Volume of H2 produced (cm3)â?? 0 0 0 0 0 0 0 5 7 14 17 18 23 24 10 13 24 30 31 45 55 15 16 27 41 45 59 69 20 17 28.5 41 48 61 71 25 18 29.5 41 48.5 61 71 30 18.5 29.5 41 48.5 61 71 1Refer to Fig 1 for graphical representation of this table (Concn./Time). Initial rates of reactions were calculated, values obtained by measuring gradients at the start of reactions; determined by drawing tangents through the first part of the graphs from the origin. Calculation to obtain gradients (G): Y = G G = Initial rate of reaction X Table 2. Initial rates for n* moles of Mg + HCl (2M) (including Y and X valuesobtained from Fig 1) [Mg] (moles)â?? Y(cm3)â?? X(sec)â?? Initial Rate Y/X (cm3/s)â?? Initial Rate (mol dm-3 /s) Â? 8.33×10-4 7 5 1.4 5.8×10-5 1.25×10-3 11.5 5 2.3 9.58×10-5 1.66×10-3 17.5 5 3.5 1.45×10-4 2.1×10-3 â?² 20 5 4 1.67×10-4 2.5×10-3 24 5 4.8 2×10-4 2.92×10-3 29 5 5.8 2.4×10-4 Refer to Fig 2 (Rate/[Mg]) The reaction order with respect to Mg was determined to be a first order reaction (Fig 2). Experiment 2. [HCl] + Mg (2.1×10-3). 2.1×10-mmoles of Mg was the desired experimental amount to use in experiment 2. Table 3. Volume of H2 produced from n Molarity of HCl + Mg (2.1×10-3) over 30 seconds [HCl] (M) â?? 2.0 1.6 1.2 0.8 0.4 Time(sec)â?? Volume of H2 produced (cm3)â?? 0 0 0 0 0 0 5 18 11 6 2 1 10 32 22 11 3 1.5 15 45 31 17 4 1.75 20 48 40 24 6 1.85 25 48.5 46 30 9 2 30 48.5 47 35 11 2 Refer to Fig 3 (Concn./Time) Initial rates were calculated from Fig 3 (Table 4). Values were plotted against concentration (Fig 4). The reaction order with respect to [HCl] was determined to be a second order reaction. This was confirmed by calculating the square of the concentration and plotting Rate/ [HCl]2 (tabulated results in Table 4). Table 4. Initial rates of reactions for n Molarity HCl + Mg (2.1×10-3) (including Y and X valuesobtained from Fig 3) and [HCl]2 values. [HCl] (M)â?? [HCl]2(M) Y(cm3)â?? X(sec)â?? Initial Rate Y/X (cm3/s)â?? Initial Rate (mol dm-3/s)Â? 2.0 4 37 10 3.7 1.54×10-4 1.6 2.56 24 10 2.4 1×10-4 1.2 1.44 12 10 1.2 5×10-5 0.8 0.64 8 10 0.8 3.33×10-5 0.4 0.16 2.5 10 0.25 1.04×10-5 Graphical representation of these results is shown in Fig?s 4 (Rate/[HCl]) and 5 (Rate/[HCl]2). Discussion In both experiments the change in concentration affected the rate positively, increases were evident. The reason for this positive change can be explained by the collision theory of reactivity. Increasing concentration of Mg/HCl, increased the initial rate, as more molecules were present increasing the probability of a collision between them. When Mg and HCl molecules collide with energy greater than a certain critical value (Wright et al., 1999), bonds are broken, leaving fragments of hydrogen, chlorine and magnesium. Fragments then combine forming new molecules, or products of the reaction. Fragments can also re-combine to become the same molecule as before; leading to no reaction. It was established that an increase in rate occurred; this increase was then assessed in greater detail to determine the mathematical relationship, between rate and concentration. The orders of the reactions were determined graphically, the order of a reaction is the power to which we have to raise the concentration to fit the rate equation (Johnson et al., 1996). Fig 2 shows the effect increasing concentrations of Mg, had on the rate of its reaction with HCl. Fig 4 shows the effect increasing concentrations of HCl, had on the rate of its reaction with Mg. The orders of reactions with respect to Mg and HCl were determined from Fig?s 2, 4 & 5. The rate equation for the whole reaction could then be determined using these orders. Fig 2 shows a linear relationship through the origin, this indicates a first order reaction with respect to [Mg]. Rate = k [Mg]1 From this is can be assumed that as the concentration of Mg doubles in a 2M HCl solution, the rate of the reaction doubles also. Fig 4 expresses a curve; this indicates a second order reaction with respect to [HCl]. Rate = k [HCl]2 Therefore it was found that as the concentration of HCl doubles when reacted with 2.1×10-3 moles Mg the rate quadruples. A square relationship was indicated, this was confirmed by plotting Rate/[HCl]2 (Fig 5), showing a straight line through the origin. From the orders the rate equation was determined. Rate = k[Mg]1[HCl]2 = k[Mg] [HCl]2 The overal order of the reaction was calculated by adding the powers of the rate equation, e.g. 3rd order (1 + 2). The rate constant (k) in the equation can also be obtained from the data, by determination of the gradient in Fig 2, and using: k = gradient [HCl]2 Conclusion Determination of reaction orders allows a more indepth understanding of a rate change, this has been achieved in this experiemnt. Further progression with the rate equation would also allow determination of the reaction mechanism, a lot can be achieved from this type of analysis. The results obtained from this investigation were acceptable, and allowed furthur analysis. I would conclude this experiment successful. Errors were kept low, however always occur to a certain degree, experimental error will have been the reason for a few rogue results, resulting in lines of best fit on graphs. Further work could involve the analysis of sulphuric acid H2SO4, as this is a dibasic acid, a different effect would occur.

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