hydrochloric acid and magnesium experiment hypothesis

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Rate of Reaction of Magnesium and Hydrochloric Acid

Introduction.

In acid-base chemical reactions, there are four main variables, which influence the rate of reaction. These include the presence or absence of catalyst, temperature, concentration, and surface area of reactants. Temperature influences the rates of reaction through kinetic energy, such that high temperatures increase the kinetic energy of reacting molecules therefore causing frequent collisions, which form products faster. High concentrations imply that more reacting molecules are at high proximity to each other therefore intermolecular collisions are frequent therefore forming products frequently. Reactants with high surface area provide a greater binding surface for other reacting molecules, and therefore increase the number of successful collisions at any moment.

To measure, the effect of each of above factors, one has to hold some factors constant during rate reaction experimentation. Therefore, this study intends to investigate the effect of concentration and surface area of reactants on the rate of chemical reactions.Magnesium metal (in form of a ribbon or powder) reacts with acids rapidly than water liberating hydrogen gas. For stance, magnesium metal reacts with hydrochloric to form magnesium chloride salt while displacing hydrogen from the acid as hydrogen gas. This is as shown in the equation below: 2HCl (aq) + Mg (s) => MgCl2 (aq) + H2 (g) Research Question: If magnesium ribbon is replaced with an equivalent weight of powered magnesium, does the rate of reaction between magnesium and hydrochloric acid double?

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Aims and objectives of the experiment

The aim of this experiment is to verify the effects of surface area of solid reactants and concentration of aqueous reactants on the rates of acid-base chemical reactions. Therefore, we sought to test the duration of reaction of equal lengths Magnesium ribbons with reducing concentrations of hydrochloric acid.

Similarly, the duration of reaction will be determined using equivalent weights of powdered Magnesium metal. The experiment will be carried at a room temperature 25 0C.

Study Variables

• The Dependent Variables: The measured duration of acid-metal reaction in seconds and the rate of gas bubbles.
• The Independent Variables:  The concentration of hydrochloric acid used and the surface area of the Magnesium metal used.

For the Magnesium ribbon, the lengths of the Magnesium ribbon used will be constant, while quantities of powdered Magnesium metal (in grams) will be equivalent to the weight of the length of magnesium ribbon used. The experiment will be carried out at room temperature (25 0C)

The study variables are summarized in the table below:

Table 1 A table of study variables and operationalization of the study variables Prediction Given that, powdered Magnesium metal has a high surface area than equivalent lengths of Magnesium ribbon, we predict that the former will have shorter duration of reaction with hydrochloric acid than the latter. We also predict that reaction of powdered Magnesium metal with highest concentration of hydrochloric acid will take the shortest duration of reaction. Hypothesis: Powdered Magnesium metal will reduce the reaction duration by a half if used in place of equivalent length of magnesium ribbon, when reacted with hydrochloric acid.

Equipment and Materials

Chemicals and Reagents The following chemicals and reagents were required in the experimentation:

• Magnesium ribbon
• Powdered Magnesium metal

0 M, 1.5M, 1.0 M and 0.5M )

• Distilled water

Apparatus and personal protection equipment

• 10 Conical flasks (100 cm 3 )
• 3 Measuring cylinders (100 cm 3 )
• Clamp stand
• Glass trough
• A Stopwatch
• Safety goggles
• Laboratory dust coat

Experimentation procedures

The experiment procedure was divided into two related investigations involving equal lengths of Magnesium ribbons and equal amounts of powdered Magnesium metal.

• First, repair your working bench by simply removing unnecessary materials. Make sure you put on your personal protective clothing and safety goggles.
• Clean the Magnesium ribbon using a sand paper to remove oxides coating its surface. This will reduce reaction errors related to impurities.
• Cut 5 equal sizes (10 cm) pieces of Magnesium from the fleshly cleaned Magnesium ribbon, weigh each of them using a digital weighing balance and record their weights.
• Wrap the magnesium pieces immediately in an aluminum foil to prevent them from being re-oxidized.
• Measure 40 ml of 3M HCl using a clean dry measuring cylinder and pour into a clean 100 ml conical flask.
• Add 40 ml of distilled water and label the conical flask with the concentration of the HCL poured.
• Repeat step 5 and 6 for 2M, 1.5M, 1M, and 0.5M HCL and keep all the acids ready on the working bench.
• Reset your stopwatch timer and prepare a gas delivery system including water bath as shown figure one below.
• Pick one piece of Magnesium ribbon drop in the first prepared acid in the conical flask and immediately start your stopwatch.
• Immediately cork the flask to the prepared gas delivery system.
• Monitor the reaction progress closely and stop your running stopwatch when the Magnesium ribbon completely dissolves in the acid and record the reaction duration in seconds in a data sheet.
• Reset your stopwatch, a repeat steps 9, 10 and 11 for the subsequent acids.
• Discard all the chemicals, wash, and rinse the conical flasks ready for another procedure.
• Repeat steps 5, 6 and 7 above.
• Add up the weights of the five 10 cm-long magnesium ribbons and obtain the average weight in grams
• Use the average weight as obtained in 15 above and weigh of an equivalent weight of Magnesium powder (for this case 0.102 grams) and pour into the first conical flask containing the 3 M HCl acid, start your stopwatch, and immediately cork the flask to the gas delivery system.
• Monitors the reaction progress and stop the stopwatch when the Magnesium powder dissolves completely in the acid.

5M HCL and clearly label your results.

• Clear your working bench.

Results and Observations

During the reaction, the water bath in the gas delivery system showed gas bubbles ascending to the gas cylinder.

At higher acid concentration, the rates of bubble forming were rapid than those in lower acid concentrations were. The most rapid gas bubbles were observed in the acid reactions with powdered Magnesium metal. The duration of reactions were recorded as shown in tables 2 and 3 below.

Table 2. A table of results showing HCl-Magnesium ribbon reaction duration (seconds) in reducing concentration

Table 3. A table of results showing HCl-Magnesium powder reaction duration (seconds) in reducing concentration

Processing and Presenting Data

Importantly, suitable acid-base indicators can be used to detect the end of the reaction accurately.

Retrieved March 8, 2012, from newton.dep.anl.gov: http://www.newton.dep.

anl.gov/askasci/chem00/

• chem00021.htm

Retrieved March 8, 2012, from chemguide.co.uk: http://www.chemguide.co.uk/physical/basicrates/surfacearea.

• Gallagher, R., & Ingram, P. (2001). Chemistry for higher tier: New coordinated science. New York: Oxford University Press. p137

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Rate and mechanism of the reaction of Mg and HCl

In the reaction between hydrochloric acid and Magnesium the overall reaction is:

$$\ce{Mg + 2HCl -> MgCl2 + H2}$$

My proposed mechanism of reaction is that $\ce{HCl}$ dissociates in $\ce{H2O}$ like so: $$\ce{H2O + HCl <=> H3O+ + Cl- }$$

And then the $\ce{H3O+}$ oxidises the magnesium like so:

$$\ce{2H3O+ + Mg -> Mg^{2+} + H2O + H2}$$

This, in an experiment to determine the rate of the reaction, measuring the $\ce{H2}$ produced would measure this mechanism, because it is the step in which hydrogen gas is evolved. Presumably these Magnesium ions are then in solution. I would therefore deduce that the step: $$\ce{2H3O+ + Mg -> Mg^{2+} + H2O + H2}$$ Is the rate-determining step, because it is the step in which magnesium ions are formed and go into solution. My rate law, deduced from these mechanisms, would look like this:

$$\ce{rate = k * [H3O+] }$$ Because it is the concentration of the $\ce{H3O+}$ which determines the formation of the Mg ions. From there, I would assume, the $\ce{Mg^{2+}}$ and the $\ce{Cl-}$ would remain in solution, as the reaction is carried out in the acid, obviously consisting of water.

From this, can I say that the rate law for the reaction

is merely: $$\ce{rate = k * [H3O+] }$$ But $\ce{[H3O+]}$ is: $$\ce{[H3O+] = (K*[H2O][HCl])/[Cl^{-}] }$$ Where K is the equilibrium constant of the acid equilibrium.

So can we write our rate law like:

$$\ce{rate = k*(K[H2O][HCl])/[Cl^{-}] }$$ Where K is the equilibrium constant and k is the rate constant and might look like this (following the Arrhenius equation): $$\ce{k = A e^{-E_a/(R T)}}$$

To make our rate law like this: $$\ce{ rate = A e^{-E_a/(R T)} *((K*[H2O][HCl])/[Cl^{-}]) }$$

My question is: Is my rate law correct for this reaction? If not, where have I gone wrong?

Additionally: This was one of our practicals. We ended up concluding that rate goes up quadratically or exponentially with concentration. This rate law would seem to indicate a linear increase, not exponential. I assume in my rate law there is no contribution from the solid Mg. Obviously in a powder, more surface area is exposed, and a high rate is experienced. Is there a way to factor this into my rate law? Thanks for the help! :)

• reaction-mechanism

2 Answers 2

Your rate law should be $$ r = k_\text{obs} \times [\ce{H3O+}]^2 = A_\text{obs} \exp\left(\frac{E_\text{a, obs}}{RT}\right) \times [\ce{H3O+}]^2\; ,$$ which nicely fits your experimental data of a quadratic dependence on the concentration for the rate.

The exponent is necessary because the reaction is seemingly second-order in the concentration of $\ce{H+}$.

• $\begingroup$ Why? What steps/mechanisms would be in the reaction if this was the rate law? :) Thanks $\endgroup$ –  Swedish Architect Commented Mar 7, 2014 at 17:16
• $\begingroup$ Is H3O order 2 because there are 2 molecules needed to oxidise every atom of Mg?? Thanks :) $\endgroup$ –  Swedish Architect Commented Mar 10, 2014 at 17:13
• 1 $\begingroup$ Well, according to the mechanism you propose: yes. If it really happens concertedly or via an intermediary step ($\ce{H+ + Mg -> Mg+ + H^{.}}$) I don't know. $\endgroup$ –  tschoppi Commented Mar 10, 2014 at 19:58
• $\begingroup$ OK. Thanks for clarifying. I would think that it would happen with an intermediate, so Mg is first oxidized to Mg+ then to Mg+2 a second time. But I don;t know. :) $\endgroup$ –  Swedish Architect Commented Mar 10, 2014 at 20:36
• 1 $\begingroup$ Yeah, it seems more reasonable. But hey, if it fits your data (and it's not valuable, to-be-published research), go ahead and propose the concerted mechanism ;) $\endgroup$ –  tschoppi Commented Mar 10, 2014 at 20:56

For the reaction $\ce{2H3O+ + Mg -> H2 + Mg^{-2}}$; the actual rate law is $r=k \times [\ce{H3O+}]^2 \times (\ce{Mg})^1.$

The most obvious error is that HCl is a strong acid, thus not in equilibrium. Cl- is merely a spectator ion.

A more accurate reaction step method is:

$\ce{H3O+ + Mg -> H2O + MgH}$

$\ce{H3O+ + MgH -> H2 + Mg^{-2}}$

The surface area of magnesium directly affects the k value, but magnesium is still a valid part of the reaction. The magnesium term may appear zeroth order if the magnesium is in far excess, so significantly lowering the initial input of magnesium will probably result in the correct rate law.

• $\begingroup$ I hope while editing I haven't messed up anything!! Feel free to rollback if i have messed up stuff :-) $\endgroup$ –  Freddy Commented Oct 30, 2014 at 7:28
• 2 $\begingroup$ I didn't think solids could be considered in a rate law - [] indicates concentration, which can only happen in liquids/gases, not solid. Clearly the surface area of the Mg has an effect - typically taken into account in the A term of the Arrhenius equation $\endgroup$ –  Swedish Architect Commented Oct 31, 2014 at 16:43

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What's the balanced equation for magnesium and hydrochloric acid?

You can watch this reaction in the video below.

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Impact of this question

Rates of reaction

Required practical 5, core practicals.

Aims of Experiment

How does the concentration of an acid affect the rate of reaction?

In this experiment you will:

• react magnesium ribbon with different concentrations of hydrochloric acid
• measure the volume of gas produced for each concentration
• use your results to work out how the rate of reaction is affected by the concentration of the acid

How does the concentration of sodium thiosulphate affect the rate of reaction?

• react different concentrations of sodium thiosulfate with hydrochloric acid
• use a stop clock to time how long it takes for the mixture to become cloudy for each concentration
• use your results to work out how the rate of reaction changes as the concentration of the sodium thiosulfate changes

Risk Asessment

As a general rule, eye protection (goggles) must be worn for all practicals.

This risk assessment is provided as an example only, and you must perform your own risk assessment before doing this experiment.

Each group will need:

magnesium strips hydrochloric acid (3 concentrations) 250 ml conical flask 100 ml gas syringe

water bath sodium thiosulfate 50 ml measuring cylinder stop clock or stopwatch 10 ml measuring cylinder

Experiment Set-up

• use a measuring cylinder to add 50 ml of 0.5 mol/dm 3 hydrochloric acid to a conical flask
• add a single 3 cm strip of magnesium ribbon to the flask, and immediately connect the gas syringe and start a timer
• at every 20 seconds, record how much gas has been produced
• when the reaction is complete, clean the apparatus as instructed by your teacher
• repeat steps 1-4 with different concentrations (1.0 mol/dm 3 , and 1.5 mol/dm 3 ) of hydrochloric acid
• use a measuring cylinder to add 10 ml of sodium thiosulfate solution to a conical flask, then add 40 ml of water (concentration 8 g/dm 3 )
• measure and record the temperature of the solution
• place the conical flask on a piece of paper with a black cross drawn on it
• use another measuring cylinder to add 10 ml of hydrochloric acid to the flask, and immediately start a timer
• when the cross is no longer visible record the time taken, and then clean the apparatus as instructed by your teacher
• 20 ml sodium thiosulfate + 30 ml water (concentration 16 g/dm 3 )
• 30 ml sodium thiosulfate + 20 ml water (concentration 24 g/dm 3 )
• 40 ml sodium thiosulfate + 10 ml water (concentration 32 g/dm 3 )
• 50 ml sodium thiosulfate + no water (concentration 40 g/dm 3 ).

Results and Analysis

For each concentration plot a graph on the same set of axes to show:

• volume of gas (ml) on the Y axis (vertical)
• time (s) on the X axis (horizontal)
• a curve of best fit

Use your graph to compare the rates of reaction with different concentrations of hydrochloric acid with magnesium. Use collision theory to explain your findings.

Plot a graph to show:

• mean time (s) on the Y axis (vertical)
• concentration (g/dm 3 ) on the X axis (horizontal)

Describe the relationship between the independent variable and the dependent variable? What were your control variables? Evaluate the two methods that you have used to investigate the effect of concentration on rate of reaction.

Exam Question and Model Answer

A chemical company makes calcium chloride by reacting calcium carbonate and hydrochloric acid. They think they can increase the rate of reaction by increasing the concentration of the acid. Describe an experiment they could do in a laboratory to be able to test this idea.

Level 1 (1-2 marks)

Add calcium carbonate powder to a conical flask. Pour hydrochloric acid into the flask, mix, and immediately attach a gas syringe. Measure how much gas has been produced every 20 seconds, and record in a table. Compare the results to find out which reaction has a faster rate.

Level 2 (3-4 marks)

Add calcium carbonate powder to a conical flask. Pour hydrochloric acid into the flask, mix, and immediately attach a gas syringe. Measure how much gas has been produced every 20 seconds, and record in a table. Repeat these steps with other concentrations of hydrochloric acid as well (keep the volume the same). Compare the results to find out which reaction has a faster rate.

Level 3 (5-6 marks)

Add 5 g of calcium carbonate powder to a conical flask. Using a measuring cylinder, pour and mix 20 ml of 0.5 mol/dm 3 hydrochloric acid into the flask, and immediately attach a gas syringe. Measure how much gas has been produced every 20 seconds, and record in a table. Repeat these steps with 1.0 mol/dm 3 and 1.5 mol/dm 3 hydrochloric acid as well (keeping the temperature, volume of acid the same, and the mass of the calcium carbonate the same). Calculate the rate of each reaction (volume ÷ time), and compare the results to find out which reaction has a faster rate.

Stoichiometry of the Reaction of Magnesium with Hydrochloric Acid

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Magnesium and HCl Lab Report

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Effects of HCl Concentration on Magnesium (HCl + Mg)

• Thread starter Turkish
• Start date Mar 26, 2007
• Tags Concentration Effects Hcl Magnesium
• Mar 26, 2007
• Floquet engineering tunes ultracold molecule interactions and produces two-axis twisting dynamics
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A PF Mountain

What do you think is going to happen, even if you cannot justify it yet? The reaction between Magnesium metal and Hydrochloric acid is, Mg (s) + 2 HCl (aq) --> H2 (g) + MgCl2 (aq) When does this reaction occur? Can the reactants just be located anywhere and this reaction will proceed, or do they need to be combined for something to happen? How do they ‘combine’ (on a molecular level) when you drop the Mg into the HCl solution?  

Yes, the more concentrated HCl solution should react with the Magnesium at a higher rate. In addition to the questions I asked above, Think how a concentrated solution of HCl differs from a dilute solution of HCl in terms of the number of H+ ions available to react with the Mg.  

Right so, increasing the concentration of the HCl (aq) solution will increase the H+ ions available..:S But my main query is.. Does the HCl split up to form H+ and Cl- ions which then react with the Mg to form MgCl2, if so what causes this separation to occur.. Is it simply becuase 'Mg wants two electrons' or something else.. Sorry I'm just not understanding how initially the reaction begins...  

Hydrochloric acid is a strong electrolyte; it will split up completely into H+ and Cl- ions in solution. The Chloride ions (Cl-) are merely spectator ions in this reaction, it is the H+ ions which are reacting with the Magnesium metal. The ionic form of the reaction is: Mg (s) + 2 H+ (aq) + 2 Cl- (aq) --> H2 (g) + Mg+2 (aq) + 2 Cl- (aq) Canceling out the Chloride ions (since they are spectators) gives us this net ionic reaction, Mg (s) + 2 H+ (aq) --> H2 (g) + Mg+2 (aq)  

• Apr 3, 2007

I tried this reaction at school becasue we were testing out the new pH probes that we got. We had HCl up to 6M and we used varying amounts, however we did not see a distinct change. But that's not to say that it should show a linear change, becasue it should. Chances are we just got some bad acid.  

There are some good reasons that you would not see large differences. First, the reaction is between a solid and a liquid. The reaction happens at the interface. The reaction rate is therefore dependent on surface area of the magnesium ribbon which is steadily growing smaller. The reaction generates bubbles on the surface that grow in size and effectively shield a significant area of the magnesium from the solution before detatching. Finally, unless the reaction is very strongly stirred, you are relying on the fairly weak currents generated by the upwelling bubbles to deliver fresh HCl solution to the surface of the Mg ribbon.  

• Feb 27, 2011

Would the HCl be considered a catalyst in this, because the more Hydrochloric acid you add, the quicker the reaction, and catalyst's are meant to speed up reactions. So to repeat, would the Hydrochloric acid be a catalyst in mixing HCl + Mg?  

A PF Planet

You are aware of the fact thread is 3 years old? It would be better to start a new thread, than to revive an old one. We don't like necromancers here. What is a complete definition of catalyst?  

A Catalyst is a substance that initiates or accelerates a chemical reaction without itself being affected. (according to http://wordnetweb.princeton.edu/perl/webwn?s=catalyst )  

Does HCl meet both conditions mentioned in the definition?  

from what i can tell, yes it does, as the HCl is not changed at the end of the reaction and it seems to accelerate the reaction, so it would be fair on my part to conclude that YES, HCl is a catalyst when mixed with Mg.  

HSStudent2014 said: the HCl is not changed at the end of the reaction

• Feb 28, 2011

MG+2HCl -> MGCl2 + H2... Ohhh.. Thanks Borek. Because it IS changed by the end, it has broken up. Therefore, it is NOT a catalyst, as it does not match the definition of a catalyst.. Unless my chemical equation is wrong (which is completely possible), in which case, I would need the correct equation in order to determine if the HCl is a catalyst or not.  

Thanks Borek :D  

FAQ: Effects of HCl Concentration on Magnesium (HCl + Mg)

What are the effects of hcl concentration on magnesium (hcl + mg) reaction.

The concentration of hydrochloric acid (HCl) in a reaction with magnesium (Mg) can significantly influence the rate and outcome of the reaction. Here are the key effects of HCl concentration on the reaction with magnesium:

1. Reaction Rate:

Higher HCl concentrations typically result in faster reaction rates with magnesium. This is because there are more HCl molecules available to collide with and react with the magnesium surface. As a result, the reaction proceeds more quickly when the HCl concentration is increased.

2. Amount of Hydrogen Gas Produced:

The concentration of HCl also affects the amount of hydrogen gas (H 2 ) produced during the reaction. With higher HCl concentrations, more H 2 gas is generated because there are more HCl molecules available to react with magnesium and release hydrogen gas as a product.

3. Temperature Changes:

The reaction between HCl and magnesium is exothermic, meaning it releases heat. Higher HCl concentrations result in more heat being generated during the reaction. This can cause the reaction mixture to become warmer or even hot, especially when concentrated HCl is used. It's important to handle concentrated acids with care due to the potential for temperature increases.

4. Reaction Completeness:

Increasing the HCl concentration can also affect the completeness of the reaction. With higher concentrations, the reaction is more likely to go to completion, meaning that a larger portion of the magnesium is consumed, and more hydrogen gas is produced before the reaction stops.

How Can HCl Concentration be Adjusted for Experiments?

To control the HCl concentration in experiments involving the reaction with magnesium, you can dilute concentrated HCl solutions with water to achieve the desired concentration. Be sure to follow safety precautions when handling concentrated acids and use proper laboratory techniques to accurately measure and mix solutions.

What Are Some Practical Applications of Understanding These Effects?

Understanding the effects of HCl concentration on the reaction with magnesium is important in various scientific and industrial applications. Some practical applications include:

1. Chemical Kinetics Studies:

Scientists use these reactions to study reaction kinetics and determine reaction mechanisms. By varying HCl concentrations, they can investigate how reaction rates change with different reactant concentrations.

2. Hydrogen Production:

The reaction between HCl and magnesium is used in laboratories and industry to produce hydrogen gas, which has applications in fuel cells, as a reducing agent, and in various chemical processes.

3. Education:

Understanding these reactions is fundamental in chemistry education and laboratory experiments, helping students learn about reaction kinetics and stoichiometry.

Overall, the effects of HCl concentration on the reaction with magnesium have implications in both academic and practical contexts, contributing to our understanding of chemical reactions and their applications.

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The effect of temperature on the rate of reaction between magnesium ribbon and hydrochloric acid

Year 11 Science Investigation (Chemistry)

The effect of temperature on the rate of reaction between magnesium ribbon and   hydrochloric acid

Aim:  The aim of this experiment is to determine whether the temperature affects the rate of reaction between magnesium ribbon and hydrochloric acid. To do this, I will cool or heat up hydrochloric acid at different temperatures (not above 60ºC) and time the reaction between magnesium ribbon and the acid. I will then calculate the rate of the reaction.    

Introduction:  In the reaction between magnesium ribbon and hydrochloric acid, the acid will react with the magnesium to produce magnesium chloride and hydrogen gas. The equation for this reaction is:    

Magnesium + Hydrochloric Acid = Magnesium Chloride + Hydrogen

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

Chemical reactions can only happen when the reactant particles (atoms, molecules, or ions) collide but not all collisions are successful in forming products. This is known as the collision theory. For the reaction to occur, the reactant particles must collide with enough energy. This energy is needed to break the bonds in the reactant particles (e.g. molecules) so that new bonds in the product molecules can be formed. This minimum amount of energy needed for the reaction to occur is called the activation energy.

Prediction:  I predict that as the temperature of hydrochloric acid increases, the rate of reaction between magnesium ribbon and acid increases. This is because increasing the temperature provides the particles with more energy to move faster (kinetic energy) and thus collide more often in a certain time. This will make it more likely that collisions result in a reaction, increasing the rate of reaction. However, since the reacting particles have more energy, more particles will have the energy to break the bonds. The collisions will be more energetic and more particles will have the activation energy needed for the reaction to occur. As there are more, effective collisions, temperature will have a large effect on the rate of reaction. Here are two diagrams to illustrate my point:      

(two diagrams)

Using the ‘Chemistry For You’ textbook (page 197 = Explaining the effect of temperature) by Lawrie Ryan I also predict that as you increase the temperature of hydrochloric acid by 10 ºC, the rate of reaction should double because the chance of particles colliding will double. This means that the temperature should be directly proportional to the rate of reaction. My hypothesis can be shown in the graph illustrated below:

Independent Variable (Variable that I am changing):  The temperature of hydrochloric acid is to be decreased to 10ºC and 20ºC by adding ice cubes to the water bath (beaker) filled with water. Yet, for the rest of the experiment, the temperature will be increased from 30ºC to 40ºC to 50ºC by adding boiling water from a kettle to the water bath. However, i f the starting temperature of the acid is more than the proposed temperature, the speed at which the acid particles collide with the magnesium ribbon will increase more. Thus, the acid particles will move with more energy increasing the rate of reaction. This means that I will have to measure the actual temperature of the acid at the beginning and end of each experiment because it is difficult to keep the temperature constant throughout the experiment.

Dependent Variable (Variable that I am measuring):  The time taken for all the magnesium ribbon to react with the hydrochloric acid is to be found with a stopwatch and rounded off to the nearest second. I will then find the rate of the reaction by dividing the time taken for the magnesium to react by 100 and rounding off the answer to 2 decimal points.

Control Variables (Variables that I am keeping constant to make it a fair test):

1. Length of Magnesium Ribbon

This is a preview of the whole essay

An increase in the length of magnesium can alter the rate of reaction because more magnesium is exposed to the acid particles. Thus, it is important to keep the length of magnesium constant. This will be decided after conducting my trials.  

2. Volume of Hydrochloric Acid  

The volume of hydrochloric acid can affect the rate of reaction. A large volume of acid would make the rate faster because there are more acid particles. I will plan to use 25cm³ of hydrochloric acid.  

3. Concentration of Hydrochloric Acid

The concentration of hydrochloric acid has already been decided for us. It will be 0.6 Molars.  

4. Surface Area of Magnesium Ribbon

The surface area of magnesium ribbon can affect the rate of reaction e.g. if the magnesium was split into two pieces, more surface area will be exposed to the acid increasing the rate of reaction.    

Apparatus List:

250ml Beaker (Water Bath)

100ml Beaker  

50ml Measuring Cylinder

Thermometer

Magnesium Ribbon (length is to be trialled)

Pair of Scissors

400ml Hydrochloric Acid (Concentration of 0.6 Molars and 25cm³ of acid for each experiment)

Hot Water/Cold Water (Volume varies for each experiment)

Ice Cubes (Varies for each experiment)  

Safety Glasses

Diagram of Apparatus Needed: (diagram)

• Set up apparatus as shown in the diagram above.
• Using a ruler and a pair of scissors, measure and cut 10 strips of magnesium ribbon required for the trial experiments (2cm x 5 and 1cm x 5). *  
• Pour some hot water or cold water in a 250ml beaker (water bath).
• Then measure 25cm ³ of hydrochloric acid (concentration of 0.6 Molars) using a 50ml measuring cylinder and add into a 100ml beaker. Place the beaker gently inside the water bath.
• Put a thermometer into the 100ml beaker to check the temperature of the acid. If the temperature has not reached the proposed temperature, add more hot water or ice cubes into the water bath until it does so. Make sure the level of the water is above the level of the acid so that the acid will fully be cooled or heated.
• When the temperature of the acid reaches the proposed temperature, put the magnesium ribbon into the acid. Start timing using a stopwatch and record the starting temperature.
• When the magnesium ribbon fully reacts with the hydrochloric acid, stop the stopwatch. Record both the temperature after the reaction and the time taken for the magnesium to completely react.  
• Do procedures 1 to 7 for temperatures 10ºC, 20ºC, 30ºC, 40ºC, and 50ºC.

Note:  In my trials I can only use a maximum of 20cm magnesium ribbon for experimentation but I will be using 15cm of magnesium altogether (2cm x 5 = 10cm and 1cm x 5=5cm, therefore 10cm + 5cm = 15cm). When it comes to doing my actual method I will use a constant length of magnesium which will be decided from the preliminary trial. The maximum length of magnesium I can use in all the actual experiments is 40cm.  

Safety: The reaction between magnesium ribbon and hydrochloric acid can be quite violent when at high temperatures and can give off acid spray so safety goggles should be worn throughout.  

Preliminary Trial:  Initially, to determine an appropriate range of temperatures and decide some suitable quantities, I conducted a preliminary investigation in which I recorded the time taken for the magnesium ribbon to react with the hydrochloric acid at various temperatures. To decrease the temperature of acid to 10ºC and 20ºC, I will add ice cubes to a 500ml beaker (water bath) filled with water. While for temperatures 30ºC, 40ºC, and 50ºC I will increase the temperature of the acid by adding hot water from a kettle to the beaker. I received the following results:

Constant Variables: 1. Volume of Hydrochloric Acid= 25cm³        

                                2. Concentration of Hydrochloric Acid = 0.6 Molars

Modification of Method: After this trial, I decided that the length of magnesium to use in my actual experiments is 1cm because changing the length to 2cm will not have much effect on the rate of reaction and that it may be wasteful if I use a larger piece. Therefore, I will change step 2 of my method to:

2. Using a ruler and a pair of scissors, measure and cut 21 strips* of 1cm magnesium ribbon.  

Each experiment will also be repeated 3 times to find an average. However, I will not plan to test how long it takes for the magnesium ribbon to react with the hydrochloric acid for 10 ºC because it takes too long. Instead, I will plan to test temperatures: 20ºC, 25ºC, 30ºC, 35ºC. 40ºC, 45ºC and 50ºC. This means that I will change step 8 to:

8. Do procedures 1-7 for temperatures 20ºC, 25ºC, 30ºC, 35ºC. 40ºC, 45ºC, and 50ºC and repeat the experiment 3 times to get an average reading to improve the reliability of my results.

I will not change anything else in my original plan because my results seem to be accurate and reliable. It can be noticed from the table that as the temperature increases, the time for the magnesium ribbon to react with hydrochloric acid increases.  

Results: Effect of temperature on the rate of reaction between magnesium ribbon and hydrochloric acid

Constant Variables:           1. Volume of Hydrochloric Acid = 25cm³              

                                          2. Concentration of Hydrochloric Acid = 0.6 Molars            

                                          3. Length of Magnesium = 1cm

                                          4. Surface Area of Magnesium = Constant as length is also constant

(Time Graph)

From the graph above I can see that an increase in temperature affects how quickly the time takes for the magnesium ribbon to disappear. As I increase the temperature of hydrochloric acid, the time taken for the reaction to occur increases. This means that hydrochloric acid with a higher temperature would make the magnesium react faster than hydrochloric acid with a lower temperature.

As the temperature increases, the time taken for the magnesium ribbon to react with the hydrochloric acid increases. Temperature affects the speed of chemical reactions when the particles collide more often in a certain time. When the hydrochloric acid particles are heated up, they have more kinetic energy causing the particles to move around and collide more quickly with the magnesium.  Therefore, the reaction time increases as the temperature increases. However, there is another reason why temperature affects the speed of the reaction. This is by making it more likely that collisions result in a reaction. This is where the temperature is increased to give the acid particles more kinetic energy and bounce off each other hard enough to start a reaction. This energy needed to start the reaction is known as the activation energy. This means that more collisions have an energy greater than the activation energy increasing the time taken for the reaction to occur.

Rates of Reaction Table:   Effect of temperature on the rate of reaction between magnesium ribbon and hydrochloric acid  

Note:  To calculate the rate of reaction I used the formula:

                               Rate of Reaction =                           100

                                      (per s)

                                                                  Time taken for Mg to react (s)

(Rate Graph)

The graph above shows that my results disagree with my prediction: that as the temperature increases by 10ºC, the rate of reaction doubles meaning that temperature is directly proportional to the rate of reaction. Thus, using the graph, the actual relationship between the temperature and the rate of reaction is that as ‘x’ (the temperature) increases so does ‘y’ (the rate of reaction) but is not directly proportional since it is a curve. This means that when the quantity on the x-axis doubles, the quantity on the y-axis more than doubles. I can show this by taking some readings from the graph:        

The total factor increase in rate is … (1 decimal point) as … + … + … + … ÷ 4 = …. This shows that 10ºC is not sufficient to double the rate of reaction because the factor increase is less than 2 i.e. … Therefore, from these results I am able to say that an increase in temperature does certainly increase the rate of reaction between magnesium ribbon and hydrochloric acid. This is because at a higher temperature the hydrochloric acid particles gain more kinetic energy making them move more faster. This rapid movement causes the hydrochloric acid particles to have more collisions with the magnesium ribbon. As the chance of collisions increases, the rate of reaction increases. However, it is not enough for hydrochloric acid particles and magnesium particles to collide. Bonds between the particles of hydrochloric acid must be broken before new particles can be made. Each acid particle needs a minimum amount of energy to break the bonds before a reaction can occur. This minimum energy is called the activation energy of the reaction. This means that when the acid particles are heated, they have more kinetic energy causing a greater proportion of them to have the required activation energy to react with the magnesium. Here is a diagram to summarise my point:

Since most of the points on the two graphs are close to the best fit curve, the results are accurate and reliable, so all my evidence supports a firm conclusion.

Evaluation:   Although I think that my experiment was sound overall, there were few points at which the accuracy was not perfect. Referring back to the graph, all my results were reasonably accurate, however there were two anomalies. The first anomalous result ( state  temp, exp no, and time) was probably due to the fact that it was not easy to see when the magnesium ribbon had fully reacted with the hydrochloric acid. This could have caused big errors in the reliability of my results. For example, the time of the chemical reaction was measured to an accuracy of 1 second. Thus, the percentage error on measuring the reaction lasting … seconds for ..ºC is …%. This could even account for the points not exactly being on the line of best fit. This is a major source of error when doing faster reactions at higher temperatures as the time could increase. The variability of all my results was reasonably similar for each temperature e.g. …, …., and …. For … ºC except the temperatures that contained anomalies.

There are several reasons for any anomalous results in this experiment. The second anomalous result ( state  temp, exp no, and time) that did not fit the pattern was probably due to not measuring the volume of acid accurately. If the volume of acid is more, this will cause the time to increase and if the volume of acid is less, this will cause the time to decrease. As I measured 25cm³ of acid in a 25cm³ measuring cylinder, the percentage error is equal to 4% (1/25 x 100 = 4%). However, this is a minor source of error.

This investigation can be improved in many ways. Inorder to improve the reliability of the method, more smaller and accurate apparatus (burettes and pipettes) can be used to measure the volume of acid. I can also try to monitor and keep my control variables as constant as possible to increase the reliability of my experiment. For example, I could measure the length of magnesium as accurately as possible and avoid wetting the apparatus (50ml beaker with 25cm³ of acid) with water for each experiment to avoid minor errors in my results. The results could also be recorded more frequently (5 repeats) to increase the reliability and a bigger range of temperatures can be included with a 5ºC interval e.g. 10ºC, 15ºC, 20ºC, 25ºC, 30ºC, 35ºC, 40ºC, 45ºC and 50ºC. However, due to practical time constraints in taking the readings for my investigation, and some consequential problems relating to time extension, I could not in fact make these adjustments.    

A definite trend can be drawn from the results for this investigation but to ensure that this is a definite conclusion the ideal method for measuring accurately the rate of reaction between magnesium ribbon and hydrochloric acid with temperature is too collect the volume of hydrogen gas given off using a syringe. The volume of gas can be measured up to a certain time period and the initial rate of reaction can be calculated. This experiment could also be carried out by using data logging equipment. This is where computer based logging and monitoring equipment is set up to measure and record the actual temperature used, the volume of hydrogen gas produced, the time taken for the magnesium ribbon to react, and the rate of the reaction. However, the ideal method was not suitable for this investigation because this equipment was not available in the laboratory and it’s too expensive. More time is also needed to time the reaction in order to calculate the rate. To extend my investigation, I could investigate the effect of temperature on the rate of reaction between magnesium and sulphuric acid. Other reactive metals and acids can also be used.    

Document Details

• Word Count 2919
• Page Count 6
• Subject Science

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