Investigating the Effect of Concentration on Reaction Time

Concentration vs Rate of Reaction

Whether you are introducing collision theory or something more demanding like reaction order, the reaction between sodium thiosulfate—Na2S2O3 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 1 experiment which focused on the reaction between sodium thiosulfate and hydrochloric acid.

Na2S2O3 (aq) + HCl (aq) → 2NaCl (aq) + S (s) + H2O (l) + SO2 (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.

This reaction and the implementation of this natural clock can be seen below in a Flinn video2.

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 ().

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?
  • How should we record and organize our data?
  • 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 Na2S2O3 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 Na2S2O3 solution. Minimizing chances for experimental error was huge.

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 3 for providing the framework we used to set up and make sense of the investigation.


1 Rate of Reaction of Sodium Thiosulfate and Hydrochloric Acid. N.p.: Flinn Scientific, n.d. Pdf.

2  "Rate of Reaction of Sodium Thiosulfate and Hydrochloric Acid..."20 Dec. 2012,  . Accessed 17 Jan. 2017.

3 "NSTA Science Store: Argument-Driven Inquiry in Chemistry: Lab ...." 1 Oct. 2014, . Accessed 17 Jan. 2017.

Supporting Information: 
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Comments 11

Tracy.Schloemer's picture
Tracy.Schloemer | Tue, 01/31/2017 - 10:14

Hey Ben-

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.


Ben Meacham's picture
Ben Meacham | Tue, 01/31/2017 - 19:57

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

Kaleb Underwood's picture
Kaleb Underwood | Thu, 02/02/2017 - 06:33

Hey Ben, 

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.

Beverly Guillebeau | Mon, 02/27/2017 - 20:08

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]= akt.  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.

Kaleb Underwood's picture
Kaleb Underwood | Wed, 03/01/2017 - 13:37

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 aA --> 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. 


John Moore's picture
John Moore | Thu, 03/08/2018 - 09:27

As defined by the International Union on Pure and Applied Chemistry, reaction rate depends on stoichiometry. You can find the defnition here: 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.

Kaleb Underwood's picture
Kaleb Underwood | Thu, 03/08/2018 - 10:24


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.


Ben Meacham's picture
Ben Meacham | Tue, 02/28/2017 - 23:30

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 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 

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 )

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! 

Kaleb Underwood's picture
Kaleb Underwood | Wed, 03/01/2017 - 19:58

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

Dan Meyers's picture
Dan Meyers | Wed, 03/07/2018 - 13:46


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

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.