How Does Acid Concentration Affect the Rate of Reaction of Magnesium with Dilute Hydrochloric Acid?

-Chemistry Coursework-

PLANNING:

Aim:

The aim of this experiment is to investigate the affect that concentration has on the rate of reaction between magnesium and dilute hydrochloric acid.

Background Information:

In a chemical reaction, there are two key components. The initial components, the reactants, are known as the starting chemicals. The resulting components are the products, as they are the chemicals made during a chemical reaction.

When Magnesium (Mg) is placed into dilute Hydrochloric Acid (HCl), it reacts to form Magnesium Chloride (MgCl2) and Hydrogen gas (H2). The balanced symbol equation for this reaction is as follows:

Mg (s) + 2HCl (aq) → MgCl2 (aq) + H2 (g)

Magnesium (Mg) is in group 2 of the Periodic Table of Elements. All elements in this group have two electrons in their outermost shell. Hydrochloric acid (HCl) is an aqueous solution because it is Hydrogen Chloride, dissolved in water (H2O). Hydrochloric acid is made up of Hydrogen and Chlorine.

H2 + Cl2 → 2HCl

Hydrogen is not classified in any set group as it shares many different properties with other elements. Chlorine is in group 7 of the Periodic Table as it has 7 electrons in its outermost shell. Hydrochloric acid is referred to as an ionic compound because of the ionic bond holding the Hydrogen (H+) ions and Chlorine (Cl-) ions. An ionic bond is a chemical bond formed when one atom transfers electrons to another atom.

When Magnesium is placed in Hydrochloric acid, it displaces the Hydrogen atom to form an ionic compound called Magnesium Chloride (MgCl2).

For this reaction to occur, the bonds between the elements in a compound need to be overcome. The rate at which this takes place is referred to as the rate of reaction. The rate of reaction is defined as the rate of loss of reactants (or rate of formation of products) during a chemical change. Rate is a measure of the change that happens in a single unit of time.

There are a number of ways to alter the rate of reaction. In some cases, it is important that a reaction must be slowed down. In most cases, it is important that a reaction takes place much faster. When altering the rate of reaction, one has to understand the Particle Collision theory. Chemical reactions occur when reacting particles collide with each other with sufficient energy to react. The minimum amount of energy required to cause this reaction is called the activation energy.

To alter the rate of reaction, one could change the temperature at which the reaction takes place. Increasing the temperature would make the particles move much faster, resulting in a higher frequency of collisions and an increased rate of reaction. Also, when the particles are travelling faster, they collide with more energy and are therefore more likely to react. Making solids into smaller pieces mean that there is more area for the particles to collide with, and therefore more frequent collisions and a higher rate of reaction. The introduction of some substances can also alter the rate of reaction.

One other important method used to change the rate of reaction (and the method I will be observing), is to change the concentration of the reactants. Changing the concentration of the reactants means that there is a change in the number of particles. Increasing the number of particles means that there will be an increased number of collisions and an increased rate of reaction.

Hypothesis:

I predict that as the concentration of Hydrochloric acid is increased, the rate of reaction with Magnesium will also increase. More specifically, I predict that if the concentration is doubled, then the rate of reaction will be doubled. I base this prediction on my knowledge of the particle collision theory. If there is double the amount of reacting acid particles, then those particles will collide with the surface of the magnesium with double the frequency and subsequently double the rate of reaction.

Apparatus:

Before doing this experiment, I did a preliminary experiment so that I would have a starting point in knowing how I would go about performing the experiment. When doing the preliminary experiment I was able do decide what apparatus to use and have chosen to use the following:

300 cm3 of Hydrochloric Acid (1 M)
Twelve, 1 centimetre strips of Magnesium
180 cm3 of Water
Glass Beaker (100 cm3)
Glass Stirring Rod

Planned Method:

Having performed the preliminary work, I have a sound plan as to what ...

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Apparatus:

300 cm3 of Hydrochloric Acid (1 M)
Twelve, 1 centimetre strips of Magnesium
180 cm3 of Water
Glass Beaker (100 cm3)
Glass Stirring Rod

Planned Method:

Having performed the preliminary work, I have a sound plan as to what measurements I will use for a beneficial experiment. The knowledge gained from the experiment allowed me to make confident choices in what measurements and amounts I would do in my final experiment. In the preliminary experiment, I used 50 cm3 of Hydrochloric Acid. For this experiment, I will use 40 cm3 of Hydrochloric Acid instead, so that it will not be difficult to measure out 0.25 M intervals between the concentrations. I will be using 4 different concentrations of acid from 1.0 M, 0.75 M, 0.5 M and 0.25 M. I have chosen to use 0.25 M intervals between each solution as I feel that this would give me a wide range of results without having an impractical number of magnesium strips to cut and acid concentrations to make.

When doing the preliminary experiment, I used various lengths of magnesium ranging from 1 centimetre to 4 centimetres. I found that the using 1 centimetre strips gave relatively fast, but easily measurable reactions. While measuring the time taken for the reaction to complete, I had difficulties with the strips of magnesium floating to the top of the acid. Stirring the beaker with the glass rod kept the strips fully submerged until the reaction was complete.

Once the all the strips have been cut and the first acid concentration made, I will start the first experiment. I will have the stopwatch ready in one hand, and one strip of magnesium in the other hand. As soon as I drop the strip, and it touches the acid, I will start the clock. I will stir the acid quickly so that the strip of magnesium will be fully submersed. I will stop the clock the second the acid stops fizzing. The time, given in minutes, will be concerted into seconds and recorded. This exact same process will be repeated 2 more times, and then the next concentration will be tested in the same manner. The rate of reaction will be found by dividing 1000 by the mean of the times taken.

Fair Testing:

Whenever carrying out any experiment, it is important that a fair test was given. When doing an experiment like this one, I’ll need to make sure that every step is performed in exactly the same way. In this particular experiment, I will ensure that the length of each magnesium strip is cut to exactly one centimetre with the ruler and not one millimetre more or less. Each solution will be measured exactly to 40 cm3 using the pipette to add water or hydrochloric acid with precise accuracy. The measuring cylinder must be read at approximately 90 degrees to avoid parallax error. The same glass stirring rod will be used in every experiment. The only variable in this experiment will be the concentrations of the hydrochloric acid.

Repeats:

I will repeat the experiment three times for each concentration of acid. I will then be able calculate a mean of the results which will allow room for experimental error and provide me with a wider range of results which will ensure that I have a much more firm conclusion.

Safety:

Every experiment needs to be executed safely. When doing this experiment, I will be working with a potentially dangerous acid, even in dilute form. Safety glasses and laboratory coats will be worn at all times, without exception. If any acid is to splash or spill onto the skin, it must be washed immediately, as well as thoroughly, and the teacher must be made aware.

OBTAINING EVIDENCE:

Actual Method:

The first step was to cut the Magnesium into 1 centimetre strips. As I had to repeat the experiment for each concentration three times and had four different acid concentrations, so I had to cut twelve strips. Each strip was cut to precisely 1 centimetre in length. The ends of every strip were cut to approximately 90 degrees so that all the strips had approximately the same mass.

Next, I made the four different acid concentrations. I measured the specific amounts of water and acid (1M) using a measuring cylinder to ensure that the correct amounts were used to obtain the accurate concentrations. When I measured the concentrations, I read the measuring cylinder at approximately 90 degrees to avoid parallax error. I used the following table to make the different concentrations.

I made three beakers of each concentration of acid. For each beaker, I dropped the 1 centimetre strip of magnesium and started the stopwatch as soon as the strip touched the water. As the magnesium was fizzing, I stirred the acid quickly with the glass stirring rod. This method ensured that the magnesium did not surface and affect my results by being exposed to the outside air. As soon as the fizzing stopped, I stopped the stopwatch and recorded the time in minutes, which was later converted into seconds.

Results:

ANALYSING:

Analysis:

All of the twelve, 1 centimetre strips of magnesium were completely dissolved in the four different hydrochloric acid concentrations. The following observations were made about the results:

The results found from the 0.25 M hydrochloric acid concentration proved to be the slowest of all the concentrations. On average, it took 1201.7 seconds for one, 1 centimetre strip of magnesium to completely dissolve. For this concentration, there were two readings that were within 15 seconds of each other, 1052 seconds and 1067 seconds respectively. However, there was one result that took 1486 seconds, approximately 427 seconds faster than the other two. In 1000 seconds, this reaction would have taken place 1201.7 times, so therefore the rate of reaction: =(1000 / 1201.7) = 0.8 (s-)

The results found from the 0.5 M hydrochloric acid concentration proved to be considerably faster than those in the 0.25 M concentration did. On average, it took 416.7 seconds for one, 1 centimetre strip of magnesium to completely dissolve. In this concentration, there are 25% more reacting particles than the 0.25 M concentration, yet the reaction takes 65.3% less time to complete. For this concentration, there were two results that were within 3 seconds of each other, 435 seconds and 439 seconds respectively. However, there was one result that took 376 seconds, approximately 61 seconds faster than the other two. According to these results, my hypothesis is only partly true. While the rate of reaction is higher with double the concentration, it unfortunately has not doubled the rate of reaction. It is however, 3 times the rate of reaction. In 1000 seconds, this reaction would have taken place 2.4 times, so therefore the rate of reaction: =(1000 / 416.7) = 2.4 (s-)
The results found from the 0.75 M hydrochloric acid concentration proved to be much faster than the 0.5 M concentration. On average, it took 172.3 seconds for one, 1 centimetre strip of magnesium to completely dissolve. In this concentration, there are 25% more reacting particles than the 0.5 M concentration, yet the reaction takes 58.7% less time to complete. The fastest of these reactions took place in 159 seconds while the longest took 185 seconds, a total range of 26. According to these results, my hypothesis is only partially supported. While the rate of reaction is higher with triple the concentration, it unfortunately has not tripled the rate of reaction. It is however, 7.25 times the rate of reaction. In 1000 seconds, this reaction would have taken place 5.8 times, so therefore the rate of reaction: =(1000 / 172.3)

= 5.8 (s-)

The results found from the 1M hydrochloric acid concentration proved to be the quickest of all the concentrations. On average, it took 71.3 seconds for one, 1 centimetre strip of magnesium to completely dissolve. In this concentration, there are 25% more reacting particles than the 0.75 M concentration, yet the reaction takes 58.6% less time to complete. That means that there were more frequent collisions than in any of the other concentrations. The fastest of these reactions took place in 69 seconds while the longest took 75 seconds, a span of 6 seconds. According to these results, my hypothesis is again, only partially supported. While the rate of reaction is higher with 4 times the concentration, unfortunately the rate of reaction is not 4 times as high. It is however, 17.5 times the rate of reaction. In 1000 seconds, this reaction would have taken place 14 times, so therefore the rate of reaction: =(1000 / 71.3)

= 14.0

The overall trend in my graphs suggests that as the concentration of hydrochloric acid gets stronger, there is an increase in the rate of reaction. This is evident by looking at both graphs and observing the fact that there is an obvious correlation between the rate of reaction and the concentration of hydrochloric acid. However, when looking at the graphs, it is clear that there is a sloping curve in both graph 1 & 2. Looking at graph 2, it shows that as the concentration is doubled, the rate of reaction is tripled (indicated in red on the graph on page 10). This does not support my hypothesis. The reason for the downward slope on graph 1 and upward slope on graph 2 may be due to the fact that the reaction taking place is an exothermic reaction. This means that the reaction between magnesium and hydrochloric acid gives off heat. This heat can easily affect the rate of reaction by increasing the frequency of the acid particles colliding with the magnesium, and also increasing the energy with which the acid particles collide with the magnesium. The increase in heat will have subsequently increased the rate of reaction.

EVALUATION:

Setting up this experiment was very straight forward and I had no difficulties. Care was taken when cutting the magnesium strips and making up the concentrations of hydrochloric acid, so that all the measurements were precise. This was the most difficult step of the setup, though it was still quite simple to perform.

When looking at the results table, it is clear that there is a large variation in some of the results, specifically the 0.5 M and 0.25 M concentrations. Despite my best efforts at reducing any unwanted variables, there were however, some experimental errors that may possibly have affected some of the results. In the 0.5 M acid concentration, one result was approximately 60 seconds slower than the other two results. In the 0.25 M acid, one result was 426 more seconds than the other two. However, the graphs I created showed a smooth sweeping line from point to point, indicating that there were no extreme anomalies in the averages.

One of the experimental errors that might have occurred and affected the results may have been the fact that I did not stir the acids consistently enough. It is quite difficult to stir the acids consistently for nearly 20 – 25 minutes. One way to improve on this problem would be to use a motorised electric stirrer. This apparatus would have stirred the acid at the same speed for as long as it takes, and I would have used one had we had the facilities in our laboratory.

Another error that could have occurred could be the width and subsequently the mass of the Magnesium strips. The Magnesium used was bought in a spool and cut from a long strip/ribbon. It was assumed that the width and thickness would be the same for all of the strips; however, this may not have been the case. One way to resolve this problem would have been to measure the mass of each strip and keep them all the same. However, this would require a great deal of time to cut them all to the same mass. It would have also been too expensive to obtain balance scales that are accurate enough for that degree of precision.

The fact that there was a layer of oxide that was covering the surface of the Magnesium could have easily affected the rate of reaction. This is because for the acid to start dissolving the magnesium, it must first dissolve the oxide coating. One way to solve this problem would have been to scrape off the oxide layer with some form of abrasive paper such as emery or sandpaper.

The primary problem was that the reaction of magnesium and hydrochloric acid is an exothermic reaction. This means that it gave off heat during the experiment, which means that the temperature could not be kept at an absolute constant. One thing I could have done to reduce the effect that temperature had on the reaction would have been to use much larger volumes of acid. The heat would have dissipated throughout the acid and it would not have affected the rate of reaction as much.

Having used three magnesium strips for each acid concentration, I increased the repetition of my experiment. Having taken the average of the repeats, I attempted to reduce the effects of the anomalous results that might have occurred. However, I would have liked to have many more repeats so that I could have discarded any extreme anomalies. This would have improved the accuracy of my results greatly.

If I had more time, I would have liked to do the experiment again so that I could record a wider range of data. I would have liked to record the mass of each magnesium strip before dissolving it in the hydrochloric acid. I would have also liked to use smaller intervals between the concentrations of acid, possibly 0.1 M increments of concentrations. I could have also used four or five magnesium strips per solution. This would have produced a much wider range of data and allowed me to have a stronger conclusion.

I would have like to have used different methods of recording evidence for the rate of reaction of magnesium and hydrochloric acid. I could have performed the experiment in a conical flask that had a rubber stopper with a gas syringe inserted through the middle of the stopper. I would have dropped the magnesium strip in the flask of acid, started the clock and replace the rubber stopper at the same time. I would have then recorded the volume of the gas syringe 10 – 15 seconds to observe how quickly the rate of reaction occurs over a period of time, as well as how long it takes for the reaction to complete. This may have proved to be much more accurate in measuring the rate of reaction.

I would have liked to have observed the affects that other variables have on the rate of reaction. If I had chopped the strips, or grinded them down to small shavings, I would have been able to observe the affect that surface area has on the rate of reaction. If I had heated one concentration to several different temperatures, I could have observed the affect that temperature has on the rate of reaction. Using different substances would have given me a wider spectrum of knowledge on the rate of reaction. I could have used Magnesium, M (s), and Sulphuric Acid, H2SO4 (aq), or Hydrochloric acid, HCl (aq), with Calcium Carbonate, CaCO3 (s).