hydrochloric acid concentration experiment

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Lesson Explainer: Effects of Temperature and Concentration on Rates of Reactions Science • Third Year of Preparatory School

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In this explainer, we will learn how to describe and explain the effect temperature and concentration have on the rate of chemical reactions.

The speed at which a chemical reaction takes place is known as the rate of reaction. Usually, the rate of reaction describes how some variable changes over a certain period of time. A common way to measure the rate of a chemical reaction is to measure how the concentrations of the reactants and products change over a certain period of time.

Definition: Rate of Reaction

• The rate of reaction measures how reactant or product concentrations change per unit time.

The rate of a chemical reaction can be affected by many factors. By changing some of these factors, the rate of reaction can be increased or decreased.

The factors that affect the rate of reaction include surface area, temperature, concentration, and the addition of catalysts. We will focus on temperature and concentration.

In order for two particles to react, they must first collide. In addition, the particles must have a certain amount of energy when they collide.

Any factor that can increase the frequency of collisions, or the energy of the particles, will likely increase the rate of reaction.

Example 1: Identifying in Which Box of Particles the Number of Collisions Will Be Greatest

The boxes below represent a chemical reaction between the red and the blue particles. In which box will the number of collisions be greatest?

A chemical reaction occurs when reactants collide with each other. The greater the number of collisions that occur, the more likely the reaction to happen and the faster the rate of reaction.

There are several factors that can affect the rate of reaction. However, from the question and diagram, we can see that we are given four boxes each containing different numbers of particles. The size of the box is also the same in each case.

If the particles are moving randomly, then the more particles there are, the more collisions there are likely to be.

We can see from the diagram that box A contains the greatest number of particles. Therefore, the number of collisions is likely to be greatest in box A.

The answer is box A.

One way to increase the number of collisions is by increasing the temperature. As the temperature increases, the particles gain energy and move faster. The faster the particles move, the more likely they are to collide with each other.

In the diagram below, the larger the arrow, the faster the particle is moving. At higher temperatures, the particles have more energy and so a larger arrow.

The effect of temperature on the rate of reaction can easily be demonstrated in a laboratory experiment. In this experiment, one effervescent tablet is put into a flask that contains hot water and a second tablet is put into a different flask that contains cold water.

The tablet reacts with the water to produce carbon dioxide gas. The experimental setup is shown below.

By measuring the volume of gas produced in each experiment, the rates of reaction can be determined and compared.

The results of this experiment are shown in the graph below:

At the higher temperature, the particles have more energy and move around faster. This increases the number of collisions between particles and increases the rate of reaction.

A faster rate of reaction increases the volume of gas produced at the start of the reaction, resulting in a steeper line on the graph. However, as the mass of the tablet and volume of water remain constant, the final amount of gas produced is the same.

Example 2: Relating Temperature to the Frequency of Collisions between Molecules

The boxes below each contain an equal number of reactant molecules. The boxes are heated to different temperatures. Which box will have the greatest frequency of collisions between molecules?

In order for two reactant molecules to react, they have to collide. There are several factors that can increase the number of collisions between reactant molecules. One of these is temperature.

We are told that each box contains the same number of reactant molecules, so the frequency of collisions is not going to be affected by a different number of molecules. However, the temperature of each box is different, and so, the main effect on the frequency of collisions will be the change in temperature.

As the temperature increases, the reactant molecules gain energy and move faster. The faster the molecules are moving, the more likely they are to collide and the greater the frequency of collisions will be.

The higher the temperature, the greater the frequency of collisions between molecules. Looking at the diagram, we can see that the box with the highest temperature is box D. Therefore, the answer is box D.

Temperature is a very important factor for controlling the rate of reactions in food. Placing food in a cool place, such as a refrigerator or freezer, slows down the chemical reactions that spoil food. As a result, food can be preserved and last longer.

High temperatures are often used when cooking food. The higher temperature increases the rate of reaction and helps cook food quicker and more thoroughly.

The effect of concentration on the rate of reaction can be explained by looking at the frequency of collisions.

Consider the reaction between the purple particles A and the green particles B shown in the diagram below.

If the concentration of B is increased, then the number of particles of B present increases. This is shown in the diagram below.

An increase in the number of particles will result in an increase in the number of collisions. A greater number of collisions causes an increase in the rate of reaction.

The effect of concentration on the rate of reaction can be demonstrated using the reaction of iron wool and oxygen.

Iron wool, also known as steel wool, can be burned in the presence of oxygen. However, the speed and intensity of this reaction changes when the concentration of oxygen changes.

When burned over a Bunsen burner, the iron wool is being burned in air. Air contains 2 1 % of oxygen, a medium to low concentration. The rate of reaction is quite low, and the iron wool burns relatively slowly.

However, when burned in pure oxygen the reaction is much more rapid and intense. The concentration of pure oxygen is ∼ 1 0 0 % , much greater than air. The increase in oxygen concentration increases the rate of reaction and results in a more vigorous and fast reaction.

These two experiments are shown in the image below.

Example 3: Explaining the Different Rates of Combustion in Air Compared with Pure Oxygen

Why is the combustion of aluminum in air slower than in pure oxygen?

• The temperature of oxygen in air is greater than in pure oxygen.
• The temperature of pure oxygen is greater than air.
• The concentration of oxygen in air is less than in pure oxygen.
• The concentration of oxygen in air is greater than in pure oxygen.

The process of combustion usually refers to the reaction of a substance with oxygen. Here, aluminum is reacted with oxygen under two different conditions.

The combustion of aluminum in air is most likely performed using a Bunsen burner. Air usually contains around 2 1 % oxygen, a relatively low amount of oxygen.

The combustion of aluminum with pure oxygen most likely involves conditions where there is ∼ 1 0 0 % oxygen. We can see that the difference between burning in air and in pure oxygen is the amount, or concentration, of oxygen present.

From this, we can conclude that the difference in the rate of combustion is because of the different concentrations of oxygen. Our answer is therefore likely to be either C or D.

Concentration can affect the rate of reaction by changing the number of reactant molecules present. The more reactant molecules there are, the greater the number of collisions that will occur between them and the faster the rate of reaction is.

As concentration increases, the rate of reaction increases.

The combustion of aluminum in air is slower because the concentration of oxygen is lower than in pure oxygen. This statement matches with choice C, and so our answer is C.

Another experiment that shows the effect of concentration on the rate of reaction is the reaction of magnesium with hydrochloric acid.

In this experiment, one conical flask contains dilute hydrochloric acid and a different flask contains concentrated hydrochloric acid. Into each conical flask is placed an identical piece of magnesium of the same size and mass.

The chemical equation for the reaction between magnesium and hydrochloric acid is M g ( ) + 2 H C l ( ) M g C l ( ) + H ( ) s a q a q g 2 2

Therefore, by measuring the volume of hydrogen gas produced over time, any change in the rate of reaction can be determined.

The setup of this experiment is shown in the image below:

By plotting a graph of the volume of hydrogen gas produced against time, the rates of reaction for each experiment can be determined. A graph showing the rate of reaction for dilute and concentrated hydrochloric acid is shown below:

The graph shows that a greater volume of hydrogen gas is produced over a short period of time when concentrated hydrochloric acid is used. This shows that the rate of reaction increases as the concentration increases.

As the concentration of hydrochloric acid increases, the number of acid particles present increases. As a result, there is a greater number of collisions between the acid and the magnesium particles, and so, there is an increase in the rate of reaction.

Example 4: Ordering Experiments with Differing Concentration by Their Rate of Reaction

A chemist performs a series of experiments to determine the effect of concentration on the rate of a reaction. They pour an equal amount of hydrochloric acid of different concentrations into four test tubes, then they place an identical piece of magnesium ribbon into each of the test tubes. The experiment setup is shown below.

From slowest to quickest, what is the likely ordering of the rate of reaction for the four experiments?

There are several factors that can affect the rate of reaction. These include concentration and surface area. In the experiment, the volume of hydrochloric acid used is kept the same. An identical piece of magnesium is also used, and so, the surface area and mass are kept the same.

The only factor that is changing is the concentration of hydrochloric acid. The concentration is greatest for experiment D and lowest in experiment B.

For a reaction to occur, the reactant molecules must collide with each other. Increasing the number of collisions increases the rate of reaction.

When the concentration is increased, the number of acid particles present in the solution increases. The increased number of acid particles will result in a greater number of collisions and therefore a faster rate of reaction.

If the rate of reaction increases as the concentration increases, then the order of the rate reaction from slowest to quickest will correspond to the order from the lowest to the greatest concentration.

From slowest to quickest, the likely ordering is B, C, A, D, which corresponds to answer choice D. The correct answer is therefore D.

Example 5: Identifying Which Set of Conditions Gives the Greatest Rate of Reaction

In a series of experiments, a student changes both the concentration and the temperature. The conditions for each experiment are shown below. In which conical flask is the rate of reaction likely to be highest?

The rate of a reaction is affected by both temperature and concentration. For a reaction to occur, reactant particles must collide with each other. Any factor that increases the number of collisions is likely to increase the rate of reaction.

As the temperature increases, the particles are given more energy and can move faster. As a result, there is likely to be a greater number of collisions and a faster rate of reaction. Therefore, the rate of reaction increases as the temperature increases.

As the concentration increases, the number of reactant particles increases. With a greater number of particles present, there is likely to be a greater number of collisions and a faster rate of reaction. Therefore, the rate of reaction increases as the concentration increases.

From the two statements above, we can conclude that the rate of reaction is likely to be highest when both the temperature and the concentration are greatest.

In the diagram above, we can see that the highest temperature is 5 0 ∘ C and the highest concentration is 2 mol/L , which occurs in experiment C.

The rate of reaction is therefore likely to be highest for experiment C.

• For a chemical reaction to occur, reactant particles must collide with each other.
• Generally, as the number of collisions between reactant particles increases, the rate of reaction increases.
• When the temperature increases, the particles gain more energy and the number of collisions increases, causing the rate of reaction to increase.
• The effect of temperature on the rate of reaction can be seen experimentally by reacting effervescent tablets with water and measuring the volume of gas produced.
• Increasing the concentration increases the number of particles present. There is a greater number of collisions, and so, the rate of reaction increases.
• The combustion of substances such as iron wool in pure oxygen is faster than in air because the concentration of oxygen is lower in air.
• The effect of concentration on the rate of reaction can be seen experimentally by reacting magnesium with different concentrations of hydrochloric acid and measuring the volume of gas produced.

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• Ben Meacham's blog

Investigating the Effect of Concentration on Reaction Time

Whether you are introducing collision theory or something more demanding like reaction order, the reaction between sodium thiosulfate—Na 2 S 2 O 3 and hydrochloric acid can provide a consistent, accurate, and engaging opportunity for investigating these topics.

A few weeks ago, I was looking for a new reaction that could be used to investigate how concentration affects reaction time. In the past, I had always used traditional reactions such as magnesium and hydrochloric acid or Alka-Seltzer and hydrochloric acid. Though both served their purpose, there would always be groups that didn’t quite get data that was consistent with what I knew the relationship to be. In most cases, this was due to ambiguous and inconsistent timing methods or simply a matter of experimental error like not ensuring the magnesium stayed in the acid without floating to the top. I wanted a reaction that would be more likely to produce consistent results from group to group, easy to execute, and was a bit more exciting than waiting for magnesium or Alka-Seltzer to disappear.

Eventually, I came across a Flinn 1 experiment which focused on the reaction between sodium thiosulfate and hydrochloric acid.

Na 2 S 2 O 3 (aq) + HCl (aq) → 2NaCl (aq) + S (s) + H 2 O (l) + SO 2 (g)

What I liked most about this reaction was the easy and consistent timing mechanism it provided my students with, which could eliminate the ambiguity and differences in timing approaches that lab groups had used in the past.

Here’s how: As the reaction proceeds, one of the products is sulfur. As more sulfur gets produced, the solution becomes more and more cloudy until eventually the solution is opaque. Because of this, the moment that you can no longer see through the solution can be used as a consistent way to stop time. When I asked my students how we would all consistently decide on when the solution is opaque, many of them suggested to place some sort of object on the other side of the beaker so that we would all stop the timer when the object was no longer visible. This naturally progressed to the idea of drawing something on the beaker itself (an X on the bottom in this case) and applying the same reasoning.

reactions series of beakers with X on bottom

series of beakers after X is blocked

This reaction and the implementation of this natural clock can be seen below in a Flinn video 2 .

Even though it is just a matter of changing from visible to opaque, I noticed that the anticipation of waiting for that X to disappear had nearly all my students hovering over their beakers anxiously waiting to stop their timer. It even got to a point where different groups started to use their phones to make time lapse videos of their reaction beakers. You can see one below. As a teacher, it was fun to watch their level of excitement over something so seemingly simple.

Though I used this experiment to primarily investigate collision theory and different factors that affect the time it takes for a reaction to complete, it could easily be used to determine something more complex like reaction order ( see the entire Flinn video from which the above clip is taken ).

I also found this lab to serve as a great opportunity for my students to play a larger role in the creation of the experimental setup since there wasn’t much complexity to it. I facilitated the design of the experiment by asking my students a series of questions that were meant to feel like it was a genuine conversation happening between scientists interested in answering a question. The PowerPoint that I used to help facilitate this discussion can be found as Supporting Information at the bottom of this post if you are logged in to ChemEd X, but the general theme followed these questions:

• What is our independent variable? How should we go about changing this?
• Should the total volume of each beaker be the same or different? Why?
• What is our dependent variable?
• Are there any variables that we should control?
• How should we go about timing our reaction?
• How should we record and organize our data?
• How are we going to figure out our concentrations in terms of Molarity?
• What are we going to do with our data once we have it? Graph it?

I don’t include students in things like this often enough and it’s important that I continue to remind myself the beneficial experience this can provide for students to get a more accurate understanding of how science operates.

However you decide to do it, the general approach to this experiment goes something like this:

1) Using a Sharpie, draw a black X on the bottom (outside) of each beaker. 2) A stock solution of 0.15 M Na 2 S 2 O 3 is used to make 5 different concentrations using different amounts of distilled water, though our tap water worked just fine too. The total volume of each solution should be the same in each beaker. 3) Add 5 mL of 2 M HCl to your first beaker to start the reaction. You can give it an initial stir to uniformly distribute the HCl. The timer starts after this initial swirl. 4) While looking down at the beaker, stop the timer the moment you see the X completely disappear from sight. 5) Do this for all your samples and start analyzing your data

After everyone had finished the experiment and analyzed their results, I was thrilled to see that the data from each group produced a graph that displayed the relationship I was looking for. Not a single group had one weird outlier or a graph with seemingly random points all over the place! Some of the groups even paid close enough attention to the fact that each beaker had different levels of “opaqueness” to them. This provided a great opportunity to talk about the benefits of qualitative evidence as well. I attribute these consistent results to two primary things:

1) Consistent timing mechanism that each group can easily reproduce 2) It is almost impossible to mess up this reaction—you’re just pouring HCl into your Na 2 S 2 O 3 solution. Minimizing chances for experimental error was huge.

Effect of Concentration on Reaction Time Graph

Though I don’t always shoot for consistent data between groups when we do a lab, I knew that the arguments would vary between groups when trying to explain why their experiment displayed the relationship it did. It is the arguments I am most interested in developing after students complete their data analysis.

Students were tasked with developing their initial argument using a Claim, Evidence, Reasoning (CER) framework. Though most boards had similar claims, they often differed in what evidence they chose to present. They all had access to the same evidence and yet different groups intentionally left out certain pieces of evidence—why? Where their boards differed the most was in their reasoning, which is meant to have them justify why their evidence makes sense based on known scientific principles. I should mention that the students had not been presented anything about collision theory before this lab and yet many of them were able to come up with a valid particle-based explanation while others either circled around ambiguity, lacked detail, or simply displayed some form of misconception. The important part of this was that they tried their best, based on the models they had running around in their heads, to explain the phenomenon and knew that it was up to the scientific community (our class) to act as a filter for sorting out valid explanations from ones that either lacked detail or could not quite account for the evidence. This is the process I love doing the most.

The lab itself took about 30 mins to do but because I involved them in the experimental setup and dedicated time to construct arguments that were presented, debated, and refined, the entire process took 3 periods (1 hr each).

I want to thank Flinn for inspiring the idea for the experiment in the first place and NSTA’s book Argument-Driven Inquiry in Chemistry 3 for providing the framework we used to set up and make sense of the investigation.

Resources 1 Rate of Reaction of Sodium Thiosulfate and Hydrochloric Acid . N.p.: Flinn Scientific, n.d. Pdf . https://www.flinnsci.com/globalassets/flinn-scientific/all-free-pdfs/dc91860.pdf 2  "Rate of Reaction of Sodium Thiosulfate and Hydrochloric Acid..."20 Dec. 2012, & https://www.youtube.com/watch?v=r4IZDPpN-bk . Accessed 17 Jan. 2017. 3 "NSTA Science Store: Argument-Driven Inquiry in Chemistry: Lab ...." 1 Oct. 2014, https://www.nsta.org/store/product_detail.aspx?id=10.2505/9781938946226 . Accessed 17 Jan. 2017.

General Safety

For Laboratory Work:  Please refer to the ACS  Guidelines for Chemical Laboratory Safety in Secondary Schools (2016) .  

For Demonstrations: Please refer to the ACS Division of Chemical Education Safety Guidelines for Chemical Demonstrations .

Other Safety resources

RAMP : Recognize hazards; Assess the risks of hazards; Minimize the risks of hazards; Prepare for emergencies

Science Practice: Analyzing and Interpreting Data

Analyzing data in 9–12 builds on K–8 and progresses to introducing more detailed statistical analysis, the comparison of data sets for consistency, and the use of models to generate and analyze data.

Analyzing data in 9–12 builds on K–8 and progresses to introducing more detailed statistical analysis, the comparison of data sets for consistency, and the use of models to generate and analyze data. Analyze data using tools, technologies, and/or models (e.g., computational, mathematical) in order to make valid and reliable scientific claims or determine an optimal design solution.

Science Practice: Asking Questions and Defining Problems

Asking questions and defining problems in grades 9–12 builds from grades K–8 experiences and progresses to formulating, refining, and evaluating empirically testable questions and design problems using models and simulations.

questions that challenge the premise(s) of an argument, the interpretation of a data set, or the suitability of a design.

Scientific questions arise in a variety of ways. They can be driven by curiosity about the world (e.g., Why is the sky blue?). They can be inspired by a model’s or theory’s predictions or by attempts to extend or refine a model or theory (e.g., How does the particle model of matter explain the incompressibility of liquids?). Or they can result from the need to provide better solutions to a problem. For example, the question of why it is impossible to siphon water above a height of 32 feet led Evangelista Torricelli (17th-century inventor of the barometer) to his discoveries about the atmosphere and the identification of a vacuum.

Questions are also important in engineering. Engineers must be able to ask probing questions in order to define an engineering problem. For example, they may ask: What is the need or desire that underlies the problem? What are the criteria (specifications) for a successful solution? What are the constraints? Other questions arise when generating possible solutions: Will this solution meet the design criteria? Can two or more ideas be combined to produce a better solution?

Science Practice: Constructing Explanations and Designing Solutions

Constructing explanations and designing solutions in 9–12 builds on K–8 experiences and progresses to explanations and designs that are supported by multiple and independent student-generated sources of evidence consistent with scientific ideas, principles, and theories.

Constructing explanations and designing solutions in 9–12 builds on K–8 experiences and progresses to explanations and designs that are supported by multiple and independent student-generated sources of evidence consistent with scientific ideas, principles, and theories. Construct and revise an explanation based on valid and reliable evidence obtained from a variety of sources (including students’ own investigations, models, theories, simulations, peer review) and the assumption that theories and laws that describe the natural world operate today as they did in the past and will continue to do so in the future.

Science Practice: Developing and Using Models

Modeling in 9–12 builds on K–8 and progresses to using, synthesizing, and developing models to predict and show relationships among variables between systems and their components in the natural and designed worlds.

Modeling in 9–12 builds on K–8 and progresses to using, synthesizing, and developing models to predict and show relationships among variables between systems and their components in the natural and designed worlds. Use a model to predict the relationships between systems or between components of a system.

Science Practice: Engaging in Argument from Evidence

Science practice: obtaining, evaluating, and communicating information.

Engaging in argument from evidence in 9–12 builds on K–8 experiences and progresses to using appropriate and sufficient evidence and scientific reasoning to defend and critique claims and explanations about natural and designed worlds. Arguments may also come from current scientific or historical episodes in science.

Engaging in argument from evidence in 9–12 builds on K–8 experiences and progresses to using appropriate and sufficient evidence and scientific reasoning to defend and critique claims and explanations about natural and designed worlds. Arguments may also come from current scientific or historical episodes in science. Evaluate the claims, evidence, and reasoning behind currently accepted explanations or solutions to determine the merits of arguments.

Science Practice: Planning and Carrying out Investigations

Planning and carrying out investigations in 9-12 builds on K-8 experiences and progresses to include investigations that provide evidence for and test conceptual, mathematical, physical, and empirical models.

Planning and carrying out investigations in 9-12 builds on K-8 experiences and progresses to include investigations that provide evidence for and test conceptual, mathematical, physical, and empirical models. Plan and conduct an investigation individually and collaboratively to produce data to serve as the basis for evidence, and in the design: decide on types, how much, and accuracy of data needed to produce reliable measurements and consider limitations on the precision of the data (e.g., number of trials, cost, risk, time), and refine the design accordingly.

HS-PS1-5 Rates of Reactions

Students who demonstrate understanding can apply scientific principles and evidence to provide an explanation about the effects of changing the temperature or concentration of the reacting particles on the rate at which a reaction occurs.

*More information about all DCI for HS-PS1 can be found at  https://www.nextgenscience.org/dci-arrangement/hs-ps1-matter-and-its-interactions and further resources at https://www.nextgenscience.org .

Assessment is limited to simple reactions in which there are only two reactants; evidence from temperature, concentration, and rate data; and qualitative relationships between rate and temperature.

Emphasis is on student reasoning that focuses on the number and energy of collisions between molecules.

All comments must abide by the ChemEd X Comment Policy , are subject to review, and may be edited. Please allow one business day for your comment to be posted, if it is accepted.

Comments 11.

This is awesome. I found this lab to be very useful too, and appreciate how you've shared how it's run in your classroom.

Thanks!  Glad I came across it and was able to reflect/share.  

Bringing Kinetics to First Year Chem

I thoroughly enjoyed reading your reflections on this activity. I use a microscale version of this reaction in my AP Chemistry class and have students calculate the order of reaction with respect to thiosulfate and hydrochloric acid. It is a very reliable procedure and the students enjoy the lab for the reasons you've discussed.

I usually don't touch on kinetics in my first year course, but this year while I was teaching it I realized that the theory of kinetics (collision theory, activation energy, catalysts, decrease of rate with time) is very accessible to first year students who have a firm grasp of the particulate nature of matter. Thank you for posting how you went through this with them, I plan on giving it a shot in my chemical reactions unit that will now include basic kinetic theory.

Kinetics Question

Thank you for sharing this lab!  I am a new teacher and really appreciate such good resources.

One question: In my textbook (Chemistry by Whitten, Davis, Peck, Stanley), the integrated rate laws use ln [A]0- ln [A]= a kt.  So when working textbook problems, I've had the students use the coefficient in calculations.  However, I noticed in the AP FRQ a is not included (such as 2004B #3) and a is not included in the given equations.  I am confused on what is the correct method and how I should be teaching this.  I would appreciate any clarification.

Integrated Rate Law Question

Hi Beverly, 

I have a coppy of the 10th edition of Whitten (though I do not use it) and it does indeed use " a " for the coefficient from the balanced equation. This is new to me and I have not seen it before. However it makes sense if you look at how they set up the integration compared to other sources.

Method 1: The rate of reaction of a first order reaction A --> Products is defined as Rate = -d[A]/dt = k [A]. This assumes A has a coefficient of 1.

Method 2: The Whitten text defines the same thing, but uses the reaction a A --> Products as the model. This leads to Rate = (1/ a )(-d[A]/dt). This affects the value of k  in Rate = k [A] and the inclusion of  a  in the integrated rate law.

k  (the "rate constant") is simply a proportionality constant, it's value just depends on how you define it. If we say that  k = ak'  then if you're asked to calculate "the rate constant of the reaction" and use Method 1 exclusively then you are solving for k . If you take in to account the stoichiometry, you are solving for k' .

Given the prevelance of not including  a  I would assume that "the rate constant" is widely considered by chemists to be the value obtained via Method 1.

Now for your concerns about practice in AP Chemistry. 

This area of possible confusion have only come up twice to my knowledge. Once in 2008 #3 and once in 2016 #5. In both situations the graders accepted either value for  k . The scoring guidelines for both exams are here:

2008 Scoring Guidelines

2016 Scoring Guidelines

The forumula included on the formula chart, combined with the precedent of these two equations leads me to belive that either method will be accepted unless a more specific question were asked.

I hope this helps. 

Definition of reaction rate

As defined by the International Union on Pure and Applied Chemistry, reaction rate depends on stoichiometry. You can find the defnition here: https://goldbook.iupac.org/html/R/R05156.html . So, if the reaction is aA --> products, the rate is defined as -(1/a)(d[A]/dt). This affects the integration, and therefore the integrated rate law, just as Kaleb says.. If the stoichiometry is A --> products, then a does not appear in the integrated rate law but only because a = 1. It appears that the AP folks allowed for both of these possibilities, which seems reasonable to me.

The version with a included is more general and gives the other version when a = 1. The distinction is important when rate constants are reported in a published paper because if the stoichiometric coefficient a is not included the rate constant value will be off by a factor of a. However, the distinction seems a lot less important when students are learning this for the first time.

I did want to clarify that the version with a included is the version that most chemists who do kinetics studies would say is correct.

IUPAC To the Rescue!

Thank you for your response and link to the Gold Book! I am glad to know that the chemistry community does have a a published, accepted standard for this (and that I was incorrect in my assertion). I agree that the distinction seems less imporant for first-time students, I am curious if this is the reasoning of the AP Test Development Committee as well and am going to reach out to see.

Kinetics Response

Hey Beverly, 

I don't teach AP so I don't want to suggest a "correct method" but here's what I'm thinking based on my own limited knowledge of integrated rate laws.  

The short answer: I don't think the coefficient ( a ) is necessary.  

Why I think  a  isn't necessary: I think your answer can be found in the difference between differential rate laws and integrated rate laws--at least it helped me understand it better.  Resource  here

Differential rate laws express the rate of reaction as a function of a change in the concentration of one or more reactants over a particular period of time, they are used to describe what is happening at the molecular level during a reaction (mechanism-focused).  

On the other hand, integrated rate laws express the reaction rate as a function of the intial concentration and a measured (actual) concentration of one or more reactants after a sepcific amount of time has passed--they are used to determine the rate constant and the reaction order from experimental data.

To me, that means that because the order of a reaction is determined experimentally, they do not represent the coefficients from a balanced equation like they would for an equilibrium expression.  In other words, the expression used for a rate law generally bears no relation to the reaction equation, and must be determined experimentally (Resource  here )

I hope that helped somewhat.  There are several people on this site that would be most likely provide a much easier answer so I can reach out to others if this didn't help.  If nothing else, I got to brush up on topics I haven't dealt with for some time! 

sodium thiosulfate pentahydrate

I checked my chemical inventory and found that I only have the hydrate. Do you think it would work?

Hydrate Will Work

It will work. Just make sure you account for the added mass from water when making your solutions of desired concentration.

Disappearing X

Great minds think alike. I posted a video post about 1.5 weeks before on this same topic.

https://www.chemedx.org/blog/disappearing-x-lab

I plan on reading your post more in depth tonight during conferences if time allows. I don't do much modeling or CER although more of this may show up as we revamp our chemistry 1 curriculum to comply with our updated state science standards.

• Chemistry Articles
• Effect Of Concentration On The Rate Of Reaction Between Sodium Thiosulphate And Hydrochloric Acid

Effect of concentration on the rate of reaction between sodium thiosulphate and hydrochloric acid

In this article, we have discussed the effect of concentration on the rate of reaction between hydrochloric acid and sodium thiosulphate.

The aim of this experiment – Understanding the effect of concentration on the rate of reaction between hydrochloric acid and sodium thiosulphate.

The reaction between Sodium thiosulphate (Na 2 S 2 O 3 ) and hydrochloric acid (HCl)

To produce a colloidal solution of sulphur, where the solution obtained is translucent.

The reaction occurs as follows: Na 2 S 2 O 3 (aq) + 2HCl (aq) → 2NaCl (aq) + H 2 O(l) + SO 2 (g) + S(s)

The above reaction when written in its ionic form: S 2 O 3 -2 (aq) + 2H + (aq) → H 2 O (l) + SO 2 (g) + S(s)

As the temperature of the system increases or as the concentration of reacting species increases the rate of precipitation of sulphur also increases. As the concentration increases, molecular collisions per unit time of the reacting species increase which can result in increased chances of product formation. This results in an increase in the rate of precipitation of sulphur. Similarly, on increasing the temperature, the kinetic energy of the reacting species increases, so the number of collisions that result in the formation of products increase leading to a faster rate of reaction.

Materials required:

The apparatus and materials required for this experiment are as follows:

• Burette of volume 50 mL
• Burette stand
• Sodium thiosulphate
• 1M Hydrochloric acid

The effect of concentration on the rate of reaction:

• Take five conical flasks, rinse them with water, and label them 1, 2, 3, 4, 5.
• Add 10 mL of sodium thiosulphate solution in flask 1, 20 mL of sodium thiosulphate solution in flask 2, 30 mL of sodium thiosulphate solution in flask 3, 40 mL of sodium thiosulphate solution in flask 4, and 50 mL of sodium thiosulphate solution in flask 5.
• Add 40 mL of distilled water in the flask 1, 30 mL of distilled water in the flask 2, 20 mL of distilled water in flask 3, 10 mL of distilled water in flask 4. This is done to adjust the volume of solution in each flask to 50 mL.
• Add 1M HCl of volume 10 mL in flask 1 with the help of a burette.
• Start the stopwatch immediately.
• Take a white tile and draw a cross mark on it.
• Add half of the HCl in the flask 1 and shake it well and start the stop-watch.
• Observe the flask and start the stop-watch as soon as the cross mark becomes invisible. Record the time taken for the process.
• Repeat the experiment by adding 10 mL HCl in flask 2, 3, 4, 5 and record the time for each.

Observation and result

Precautions to be taken during the experiment:

• Thoroughly wash the apparatus.
• The solutions taken for this experiment should be measured accurately.
• Use the same tile for all the observations.
• Stay alert while you start and stop the stop-watch.

1. What is the amount of sodium thiosulphate added in flask 1?

Ans: 10 mL.

2. What is the amount of distilled water added in flask 1?

Ans: 40 mL.

3. What is the amount of HCl added in all the flasks?

4. name the two solutions used for the experiment..

Ans: Sodium thiosulphate and hydrochloric acid.

5. How many conical flasks are required for this experiment?

Put your understanding of this concept to test by answering a few MCQs. Click ‘Start Quiz’ to begin!

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• » Acid-base titration
• » Hydrochloric acid titration

Titration of hydrochloric acid with sodium hydroxide

General remarks.

Determination of hydrochloric acid concentration is probably the most often discussed example of acid-base titration. Both acid and base are strong, which not only makes determination of end point easy (steep part of the curve is long), but also means that calculation of titration curve and equivalence point are pretty straightforward.

This is a simple neutralization reaction:

HCl + NaOH → NaCl + H 2 O

It is worth of noting, that - as we can assume both acid and base to be completely dissociated - net ionic reaction is just

H + + OH - → H 2 O

which is the simplest form of neutralization reaction possible.

In the reality every acid and every base - no matter how strong - have some dissociation equilibria described by dissociation constant. In this particular case K a for HCl is listed as 10 4 (which means it can be safely neglected) and dissociation constant K b for NaOH is listed as 0.6 - which means sometimes it has to be taken into account.

sample size

Depending on the titrant concentration (0.2 M or 0.1 M), and assuming 50 mL burette, aliquot taken for titration should contain about 0.26-0.33 g (0.13-0.16 g) of hydrochloric acid (7-9 or 3.5-4.5 millimoles).

end point detection

Equivalence point of strong acid titration is usually listed as exactly 7.00. That's not necesarilly the case, as it depends on the solution temperature and ionic strength of the solution, besides, slight hydrolysis of NaOH shifts pH down by about 0.02 unit. Not that it changes much - we are still very close to 7. Thus the best indicator of those listed on pH indicators preparation page is bromothymol blue . However, as we have discussed on the acid-base titration end point detection page, unless we are dealing with a diluted solution (in the range of 0.001 M) we can use almost any indicator that gives observable color change in the pH 4-10 range. Thus we can safely use the most popular phenolphthalein and titrate to the first visible color change.

Color change of phenolphthalein during titration - on the left, colorless solution before end point, on the right - pink solution after end point. Note we have to end titration at first sight of color change, before color gets saturated.

solutions used

To perform titration we will need titrant - 0.2 M or 0.1 M sodium hydroxide solution , indicator - phenolphthalein solution and some amount of distilled water to dilute hydrochloric acid sample.

• Pipette aliquot of hydrochloric acid solution into 250mL Erlenmeyer flask.
• Dilute with distilled water to about 100 mL.
• Add 2-3 drops of phenolphthalein solution.
• Titrate with NaOH solution till the first color change.

result calculation

According to the reaction equation

Hydrochloric acid reacts with sodium hydroxide on the 1:1 basis. That makes calculation especially easy - when we calculate number of moles of NaOH used it will be already number of moles of HCl titrated.

To calculate hydrochloric acid solution concentration use EBAS - stoichiometry calculator . Download determination of hydrochloric acid concentration reaction file, open it with the free trial version of the stoichiometry calculator .

Click n=CV button above NaOH in the input frame, enter volume and concentration of the titrant used. Click Use button. Read number of moles and mass of hydrochloric acid in the titrated sample in the output frame. Click n=CV button in the output frame below hydrochloric acid, enter volume of the pipetted sample, read hydrochloric acid concentration.

sources of errors

Apart from general sources of titration errors , when titrating hydrochloric acid we should pay special attention to titrant. Sodium hydroxide solutions are not stable as they tend to absorb atmospheric carbon dioxide. Hydrochloric acid is much stronger than carbonic acid, so it will slowly expel carbon dioxide from the solution, but initially presence of carbonates will mean that to reach end point we need to add axcess of titrant.

Need more info?

Aqueous Acid-base Equilibria and Titrations

by Robert de Levie

(commissions earned)

Complete list of suggested books

last modified on October 27 2022, 21:28:27.

Dean's Analytical Chemistry Handbook - tam sa fajne (chyba) listy co czym oznaczac, warto obejrzec

Technical info: $section - alkalimetry, $subsection - , $page - HCl, $url_call - acid-base-titration-hydrochloric-acid

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The effect of temperature on reaction rate

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What effect does temperature have on reaction rates? With a little sodium thiosulfate and hydrochloric acid, students will be able to discover just that

Complete the table provided to give a clear view of the data collected, and explore temperature, reaction rates, and collision theory.

This experiment should take 60 minutes. 

• Eye protection.
• Conical flask, 250 cm 3 
• Measuring cylinder, 10 cm 3  
• Measuring cylinder, 50 cm 3  
• Sodium thiosulfate solution 40 g dm –3
• Hydrochloric acid 2 mol dm –3

Health, safety and technical notes

• Read our standard health and safety guidance .
• Wear eye protection.
• Ensure good ventilation, use fume cupboard if necessary. 
• Sulfur dioxide is formed as a by-product, see CLEAPSS Hazcard HC097 .
• For more information on sodium thiosulfate, see CLEAPSS Hazcards HC095a .
• Hydrochloric acid is an irritant, see CLEAPSS Hazcard HC047a .
• Put 10 cm 3 of sodium thiosulfate solution and 40 cm 3 of water into a conical flask. Measure 5 cm 3 of dilute hydrochloric acid in a small measuring cylinder.
• Warm the thiosulfate solution in the flask if necessary to bring it to the required temperature. The object is to repeat the experiment five times with temperatures in the range 15–55 °C.
• Put the conical flask over a piece of paper with a cross drawn on it.
• Add the acid and start the clock. Swirl the flask to mix the solutions and place it on a piece of white paper marked with a cross. Take the initial temperature of the mixture
• Look down at the cross from above. When the cross disappears, stop the clock and note the time taken. Record the final temperature of the mixture in the flask.
• As soon as possible, pour the solution down the sink (in the fume cupboard if possible) and wash away.
• Record your findings on the table provided.

The method for this experiment is best understood when the teacher demonstrates it first.

The endpoint can be measured with a light sensor connected to a data-logger. A light sensor set up as a colorimeter can be used to monitor the precipitation on a computer – clamp a light sensor against a plastic cuvette filled with the reactants.

The result, in the form of graphs on the computer, provides very useful material for analysis using data logging software.

The software shows the change on a graph, and this tends to yield more detail than the end-point approach used in this experiment.

The rate of change can be measured from the graph slope or the time taken for a change to occur. 

As soon as the reaction is complete, pour the solutions away, preferably into the fume cupboard sink. Wash away with plenty of water. This is particularly important with solutions used at higher temperatures. 

Questions 

• For each set of results, calculate the value of 1/time. (This value can be taken as a measure of the rate of reaction for this experiment).
• Plot a graph of 1/time on the vertical (y) axis and average temperature on the horizontal (x) axis.

More resources

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The effect of temperature on reaction rate - teacher notes

The effect of temperature on reaction rate - student sheet, additional information.

This practical is part of our  Classic chemistry experiments  collection.

• 11-14 years
• 14-16 years
• 16-18 years
• Practical experiments
• Physical chemistry
• Rates of reaction

Specification

• Rates of reaction can be increased: by increasing the temperature
• (a) practical methods used to determine the rate of reaction – gas collection, loss of mass and precipitation (including using data-logging apparatus)
• (b) the effect of changes in temperature, concentration (pressure) and surface area on rate of reaction
• (b) how to calculate rates from experimental data and how to establish the relationship between reactant concentrations and rate
• 2.9.1 recall how factors, including concentration, pressure, temperature and catalyst, affect the rate of a chemical reaction;
• 2.3.4 describe and explain the effects on rates of reaction when there are changes in: temperature; concentration; frequency and energy of collisions between particles; and changes in particle size in terms of surface area to volume ratio.
• 7. Investigate the effect of a number of variables on the rate of chemical reactions including the production of common gases and biochemical reactions.
• 9. Consider chemical reactions in terms of energy, using the terms exothermic, endothermic and activation energy, and use simple energy profile diagrams to illustrate energy changes.
• 2. Develop and use models to describe the nature of matter; demonstrate how they provide a simple way to to account for the conservation of mass, changes of state, physical change, chemical change, mixtures, and their separation.
• 4. Classify substances as elements, compounds, mixtures, metals, non-metals, solids, liquids, gases and solutions.
• Mandatory experiment 6.2 - Studying the effects on the reaction rate of (i) concentration and (ii) temperature, using sodium thiosulfate solution and hydrochloric acid
• Temperature.
• Activation energy and influence of temperature on the rate of reaction, using reaction profile diagrams.
• AT.5 Making and recording of appropriate observations during chemical reactions including changes in temperature and the measurement of rates of reaction by a variety of methods such as production of gas and colour change.
• Factors which affect the rates of chemical reactions include: the concentrations of reactants in solution, the pressure of reacting gases, the surface area of solid reactants, the temperature and the presence of catalysts.
• Students should be able to recall how changing these factors affects the rate of chemical reactions.
• Describe the effect of changes in temperature, concentration, pressure, and surface area on rate of reaction.
• Explain the effects on rates of reaction of changes in temperature, concentration and pressure in terms of the frequency and energy of collision between particles.
• 7.1b observing a colour change (in the reaction between sodium thiosulfate and hydrochloric acid)
• 5 Making and recording of appropriate observations during chemical reactions including changes in temperature and the measurement of rates of reaction by a variety of methods such as production of gas and colour change
• 7.4 Explain the effects on rates of reaction of changes in temperature, concentration, surface area to volume ratio of a solid and pressure (on reactions involving gases) in terms of frequency and/or energy of collisions between particles
• 8 Investigation the effect of surface area, concentration and temperature on the rate of a chemical reaction
• Making and recording of appropriate observations during chemical reactions including changes in temperature and the measurement of rates of reaction by a variety of methods such as production of gas and colour change
• C6.2.1 describe the effect on rate of reaction of changes in temperature, concentration, pressure, and surface area
• C6.2.2 explain the effects on rates of reaction of changes in temperature, concentration and pressure in terms of frequency and energy of collision between particles
• C5 Investigation the effect of surface area, concentration and temperature on the rate of a chemical reaction
• C5.1c describe the effect of changes in temperature, concentration, pressure, and surface area on rate of reaction
• C5.1d explain the effects on rates of reaction of changes in temperature, concentration and pressure in terms of frequency and energy of collision between particles
• C5.2c describe the effect of changes in temperature, concentration, pressure, and surface area on rate of reaction
• C5.2d explain the effects on rates of reaction of changes in temperature, concentration and pressure in terms of frequency and energy of collision between particles
• an increase in temperature makes molecules move faster, resulting in increased collisions and rates of reaction
• lower temperatures result in decreased collisions and rates of reaction
• temperature

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• Effect of Concentration on Rate of Reaction

Introduction

In a reaction, many elements or compounds react together to form one or more new products based on the basic elements of the compounds involved in the chemical reaction. The reaction may occur between solids, liquids, and gases, and the rate of each reaction is different from another which depends on a number of factors. Some chemical compounds called catalysts or enzymes help in accelerating the rate of reaction. This rate of reaction is also dependent on the amount or concentration of substances involved in the reaction.

In this article, the effect of concentration on the rate of the reaction between sodium thiosulphate and hydrochloric acid has been discussed in the form of an experiment.

About the Rate of Reaction

In any chemical reaction, several reactants react to form one or more new products. These reactants can be gases, solids or even liquids. The rate of reaction depends on many determinants or factors. The presence of catalysts can also accelerate the speed of any reaction. One of the essential components that determine the rate of reaction is the concentration of the reactants taking part in the chemical reaction. In this article, we are going to study the effect of concentration on reaction rate by observing a chemical reaction. The results obtained from the reaction between Na 2 S 2 O 3 and HCl helps you to understand the topic more clearly.

To understand the effect of concentration on the rate of reaction between sodium thiosulphate (Na 2 S 2 O 3 ) and hydrochloric acid (HCl).

The rate of the reaction directly depends on the products of the molar concentration of reactants. In this experiment, we will study the reaction between Sodium thiosulphate (Na 2 S 2 O 3 ) with hydrochloric acid (HCl). 

Na 2 S 2 O 3 (aq) + 2HCL (aq) → H 2 O (l) + SO 2 (g)+ 2 NaCl (aq) + S (s)

We can also write the above reaction in ionic form as:

S 2 O 3 -2 (aq) + 2H + (aq) → H 2 O (l) + S (s) + SO 2 (g)

The solution obtained after the reaction is opaque and has a milky appearance due to the presence of sulphur. If we increase the temperature or concentration of the reactants, then the rate of precipitation of sulphur also increases. It happens because when concentration increases, molecular collisions also increase per unit time, which results in a fast rate of product formation.

The reaction is between the aqueous solutions of Na 2 S 2 O 3 and HCl, the rate of which directly depends on the product of the molar concentration of each component of the reaction.

The chemical reaction formula between HCl and Na 2 S 2 O 3 is as follows:

Here, we can see the products of the reaction are water , sulphur dioxide (SO 2 ), sodium chloride (NaCl) and a sulphur atom . Where SO 2  is formed in the gaseous state, NaCl is formed in the aqueous state and sulphur collects in the solid-state.

With the increase in the temperature of the system, the precipitation of sulphur increases. With the increase in concentration, a collision between molecules also increases which increases the chances of getting the products of the reaction, and an increase in temperature provides more kinetic energy to the reaction which in turn increases the rate of reaction, thereby, resulting in faster production of products.

Materials Required

The materials and apparatus required for conducting the reaction are as follows:

Five flasks of 100ml each 

 Two burettes of volume 50 ml each

Burette stand

Sodium thiosulphate

1M Hydrochloric acid

Five conical flasks (100 ml)

Two burettes

Burette Stand

Sodium Thiosulphate

1M Hydrochloric Acid Solution

First of all, take five conical flasks and rinse them with water to clean any residue. Now, label them as 1, 2, 3, 4 and 5 respectively.

Draw the cross mark on any white tile.

Take a burette and add 10 ml of Na 2 S 2 O 3 solution in flask 1 using it. Similarly, add 20 ml of Na 2 S 2 O 3 solution in flask 2, 30 ml in flask 3, 40 ml in flask 4 and 50 ml solution in flask 5.

Now, add distilled water in every flask such that the combined volume of water and Na 2 S 2 O 3 solution becomes 50 ml. It means we need to add 40ml, 30ml, 20 ml and 10 ml distilled water in flasks 1, 2, 3 and 4 respectively.

Take 10 ml of 1M HCl solution in a test tube using the burette. Add it in flask 1, which contains 40 ml water and 10 ml Na 2 S 2 O 3 . Shake it thoroughly and then start the stopwatch immediately.

Place the flask on the white tile having a cross mark. Observe the cross mark from the top and stop the stopwatch as the cross mark becomes invisible. Note the time taken for the whole process.  

Repeat the same procedure with flask 2, 3, 4 and 5. Note the time when the cross mark becomes invisible in every container.

Wash the flasks and add 10ml of sodium thiosulphate in the first flask and add 10 ml more to each subsequent flask.

Add 40ml of distilled water in the first flask, 30ml in the second, 20ml in the third, 10ml in the fourth and none in the fifth flask.

10ml HCl is to be added to the first flask with the help of a burette and the stopwatch should be started immediately.

Take a white tile and put a cross mark on it distinctly. Put the first flask on the time and observe.

Observe till the solution is milky and the mark on the tile is visible and note the time right there.

Observations and Results

(Image will be uploaded soon)

The above image shows the graph between 1/t (on the y-axis) and the concentration of Na 2 S 2 O 3 (on the x-axis). We will obtain a straight sloping line, as shown in the figure.  

From the above graph, it is clear that 1/t is directly proportional to the concentration of Na 2 S 2 O 3 solution taken during the experiment. We know that 1/t is the direct measure of the rate of reaction. Hence, the pace of chemical reaction in this case directly depends on the concentration of Na 2 S 2 O 3 . However, it doesn't mean that the speed of any chemical reaction doesn't depend on conc. of HCl. We will see a similar result if we keep the concentration of sodium thiosulphate constant and raise the concentration of hydrochloric acid. This experiment clearly shows the effect of concentration on Rate of Reaction.

The product formed from the reaction is milky in appearance due to the presence of sulphur. Increasing the temperature increases the rate of precipitation. 

Noting the time and plotting a graph with 1/t on the y-axis and the concentration of Na 2 S 2 O 3  on the x-axis where t is the time taken to form products at different concentrations. It was observed that the graph shows a straight sloping line which means that 1/t is directly proportional to the concentration of Na 2 S 2 O 3 which means increasing the concentration increased the rate of reaction as well. 

If the concentration of sodium thiosulphate was kept constant and the concentration of hydrochloric acid was increased, the same trend would be observed. 

FAQs on Effect of Concentration on Rate of Reaction

1. What are the determining factors which affect the rate of reaction?

From the above experiment conducted, it is observed that the concentration of reactants has a direct effect on the rate of reaction. Other factors include temperature conditions given to the reaction, state of existence of the reactants, presence of catalysts and changes in the surface area affecting the rate of reaction.

2. What is the amount of HCl added to the flasks?

10 ml of HCl was added to each of the flasks.

To know more about the theory, materials required, observation, results, and conclusion of the experiment on the effect of concentration on the rate of reaction between sodium thiosulphate and hydrochloric acid, visit Vedantu's website or download the app. Vedantu gives you solutions to your queries as well as free resources to study from which you can download in PDF format and access anytime, anywhere! 

3. What was the result of the reaction?

The result of the reaction can be summed up as follows:

It was observed that in a reaction, with an increase in the concentration of sodium thiosulphate gradually while keeping the concentration of hydrochloric acid constant, the rate of reaction has increased slowly. If “t” is the amount of time taken for the products to form then 1/t is directly proportional to the concentration of sodium thiosulphate. 

The same result will be obtained if the concentration of hydrochloric acid was increased gradually while the concentration of sodium thiosulphate was kept constant.

4. What are the precautions to be taken during the experiment to study the effect of concentration on reaction rate?

First of all, one must ensure that the apparatus must be thoroughly clean before starting the process. Any impurities can lead to inaccuracy during the experiment. Moreover, it is essential to measure the volumes of the distilled water, sodium thiosulphate and hydrochloric acid accurately. You must use the same tile with the same cross mark for all the observations. The temperature variation can also affect the rate of reaction. Hence, it is essential to complete to prevent any temperature variation. Start the stopwatch immediately as you add the HCl in the solution. Finally, view the cross mark on the tile from the top from the same height during the observations.

5. What are the determinants that can affect the rate of reaction?

There is a direct effect of concentration on Rate of Reaction, as shown in the above experiment. The physical state of the reactants' surface area can also affect the pace of any reaction. For instance, if any metal reacts with gas, then only the molecules present at the surface of the metal can react with gas molecules. Hence, we can increase the surface area of reactants by cutting them in pieces to increase its rate of reaction. An increase in temperature can also enhance the rate of reaction because it will boost the kinetic energy of reactant particles. A presence of a catalyst can also accelerate the chemical reaction. 

Reaction - Magnesium & Hydrochloric Acid ( OCR A Level Chemistry )

Revision note.

PAG 9.3: Rate of Reaction - Magnesium & Hydrochloric Acid

• This reaction can be used to investigate the effect of varying the concentration of the acid while keeping the temperature constant
• However, in this case the gas hydrogen is too low in density so the mass change will be far too small to register on a laboratory balance
• The choice of the size of the gas syringe needs to be considered and the quantities of reagents judged accordingly so that a reasonable volume of gas can be evolved and also recorded

The rate of reaction between magnesium and dilute hydrochloric acid can be measured using a gas syringe and stopwatch

Steps in the procedure

• This can be done by assembling everything without the acid or magnesium and trying to move the plunger
• If you feel resistance the apparatus is gas tight
• The acid will go in first because it is quicker to drop a piece of magnesium ribbon in than to pour in the acid
• To vary the concentration of the acid you need to dilute it by measuring portions of acid in a measuring cylinder and then portions of distilled water in another measuring cylinder and adding them to the conical flask
• Choose a suitable volume of acid to match the size of the flask, e.g. 40 cm 3
• You don't want to use acid that is stronger than 2 mol dm -3  because the reaction will be too fast, so its best to start with 2 mol dm -3 and perform a serial dilution, e.g. 40 cm 3 (acid) + 0 cm 3  (water), 35 cm 3  (acid) + 5 cm 3 (water), etc

Practical tips

• Make sure the plunger is fully inserted before you start the experiment otherwise you will have a volume error
• If the magnesium does not look new and shiny, you may need to clean the surface with a bit of sandpaper
• Make sure the plunger is secure and does not fall out of the barrel if the volume exceeds 100 cm 3

Specimen Results

• Here is a set of typical results for this experiment

Rate of reaction between magnesium and acid results table

Graph showing typical results for the rate of reaction between hydrochloric acid and magnesium at different concentrations

• The lines of best fit are drawn for each concentration on the same graph
• A tangent is then drawn starting from (0,0) since this method is to find the initial rate of reaction
• The gradient of the tangent is determined which gives the rate of reaction
• In the example above, the rate of reaction for 2.0 mol dm -3 acid  is

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Author: Richard

Richard has taught Chemistry for over 15 years as well as working as a science tutor, examiner, content creator and author. He wasn’t the greatest at exams and only discovered how to revise in his final year at university. That knowledge made him want to help students learn how to revise, challenge them to think about what they actually know and hopefully succeed; so here he is, happily, at SME.

Synthesis and characterization of a new nonionic surfactant with anticorrosive activity of aluminum in 1.0 M hydrochloric acid solution

• Al Jahdaly, Badreah A.

New nonionic surfactant LMAEO-3 has been synthesized and characterized using Fourier-transform infrared spectroscopy (FT-IR) and proton nuclear magnetic resonance ( 1 HNMR). An analysis of the critical micelle concentration (CMC) ratio was conducted, which reached 11 × 10 ‑5 M. This surfactant protected aluminum from reacting with 1.0 M hydrochloric acid. In the process of checking its efficiency, electrochemical impedance spectroscopy (EIS) was employed. The double layer's resistance and capacity increased significantly due to the experiment, and it had a percentage of efficiency of 82.8%. Further, the potentiodynamic polarization study demonstrated that LMAEO-3 is a mixed-type inhibitor that inhibits current density as surfactant concentrations increase, making sense from electron frequency modulation (EFM). Observe how Langmuir's theory applies to the adsorption of this surfactant on aluminum. These results were in agreement with the theoretical results derived from Materials Studio and atomic force microscopy (AFM) testing results, which demonstrated the smoothness of the aluminum surface when only exposed to hydrochloric acid in the presence of a surfactant.

• Nonionic surfactant;
• Smoothness;
• Electrochemical investigation;
• Theoretical studies