Crystal Violet Kinetics Lab

Crystal Violet Kinetics Lab
Crystal Violet Kinetics In this Crystal Violet Kinetics Lab, many procedures were performed order
to experimentally determine the rate constant(k), for the rate of the reaction of the crystal violet
solution and sodium hydroxide. To start, a Beers Law calibration experiment was conduction. Also,
two kinetic experiments were ran using both 5ml and 10ml of sedum hydroxide in order to
determine the reaction order(k^1), of the sodium hydroxide. The Beers Law calibration experiment
used many concentrations of crystal violet solutions. Each of these solutions were test and analyzed
in order to determine the absorbance of each concentration The results were than graphed and
produced a slope of 1.00E05 with an intercept of –2.21E–02. After the
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Iodine Clock Reaction Lab Report
For this Landolt Iodine Clock reaction, the concentration of the Potassium Iodate (KIO3), the
concentration of Sodium Bisulphite (NaHSO3), and the temperature of these two solutions were
varied in separate experiments to examine the effect on the rate of the reaction. The three methods
displayed in the method section were utilised to execute these experiments.
The averaged data for varying the KIO3 concentration is displayed in Table 1 and plotted onto
Graph 1. When varying the concentration of KIO3 – 0.134, 0.100, 0.080 and 0.067M – at 30°C, it is
apparent that as the concentration decreases, the reaction rate decreases which supports the
hypothesis. This can be seen as the lowest concentration of KIO3 was 0.067M with the duration of
the reaction being 5.85s. The time decreased approximately 25% for each new KIO3 concentration
until it reached 2.95s at 0.134M.
The data for varying the NaHSO3 concentration – 0.036, 0.028, 0.023, 0.019M – at 30°C, was
averaged and illustrated in Table 3 then graphed onto Graph 3. From the table and the graph, it is
evident that as the concentration increases, the rate of reaction increases which supports the
hypothesis. The greatest concentration was at 0.036M and duration of the reaction was 5.77s which
then increased approximately 48% for each new concentration of NaHSO3 solution until the lowest
concentration, 0.019M with 18.27s. It is clear that the duration of the reaction and the percentage
increase between each concentration is

Chemical Reaction Kinetics
Reaction Kinetics
Kinetics of reactions, otherwise known as chemical kinetics, is the study of how particles and bonds
between particles change in a chemical reaction over time. These changes can be viewed at a
molecular level through the use of reaction rates. Reactions rates tell us how fast or how slow the
change is happening at this level, or how the reaction depends on the time.
In Chapter 14 of the textbook, it was given that a chemical reaction only occurs if there is a collision
between particles. The greater the amount of particles, then the greater the amoutn of collisions in a
given amount of time, the faster the rate of the chemical reation. So the rate of a reaction depends on
the concentration of the particles in the reation. ... Show more content on Helpwriting.net ...
If a chemical reaction has a slow rate, a relatively small fraction of molecules reacts to form
products in a given period of time." (Chemistry, 557) This can be applied to the rate reactions for
chemicals. How much of the reactants are used in a certain amount of time to produce the products,
is called the rate of the reaction. When finding the reaction rate of a chemical reaction, many factors
need to be included. First, we need to know what we are looking for. In Chemistry Structures and
Properties, the author provides examples such as the speed of a car being expressed in miles per
hour, or weight lost in lbs per week. We would have speed and weight, respectively, as units for our
answers. In the reaction A2 + B2 –> 2AB, for every 1 mol of A2 and B2 used, 2 mols of AB are
produced. In a chemical reaction the reactants are used up in order to form the products. In this case
the A2 and B2 are the reactants and AB is the product.
In order to find the numerical rate of a reaction, more than just the balanced equation is needed. In
the gas example above, the rate at which the gas will be consumed, and the distance, or miles, is
needed in order to find how much will be needed. This is where rate laws come into effect. Rate
laws use the balanced equation to express the relationship between the reations rate and the
reactions concentration. (Chemistry,

The Effect Of Concentration On The Rate Of Reaction By...
The hypothesis of this experiment was that an increase in concentration will increase the rate of
reaction by decreasing the time taken for the colour change to occur. Figure 1.1 shows time on the x
axis and concentration on the y axis, it can be seen that as [reactant] decreases, time increases.Thus,
the results correlate with the hypothesis that an increase in concentration will produce a faster
reaction time and subsequently a larger reaction rate – regardless of change in [KIO3] or
[Na2S2O5], both reactants produce the same result.
The second part of the hypothesis was that an increase in [B] and decrease in [A] will still increase
the overall rate of reaction, however it will cause a slow spread of colour compared to a rapid colour
change. This wasn't specifically recorded by timing, however, it was noted on the 9/5/16 that when
lower concentrations of KIO3 were tested, there was a slow spread of colour. On the 13/5/16, when
testing Na2S2O5, there was no vast change in colour spread time. Though, all the concentrations
tested of Na2S2O5 were fairly similar. If, for example, 1M Na2s2O5 was tested against 0.1M KIO3,
the reaction would definitely occur quickly, however, the spread of colour could be slow since a
large quantity of I2 has been used up to cancel the significantly larger amount of HSO3– ions. For
what has been observed, the second part of the hypothesis is plausible. Though, this is an area of
investigation which needs to be further investigated to be

Bromophenol Blue Experiment
The plot that is most linear according to the line of best fit is Figure 2, f(x) = –0.00310x –0.4269
with a R^2 value of 0.9973. Comparatively, Figure 1 (zeroth reaction order) has a best fit line of f(x)
= –0.00210(x) + 1.3426 with a R^2 value of 0.9094, and Figure 3 (second reaction order) has a best
fit line of f(x) = 0.00550x+0.2716 with a R^2 value of 0.9559. So, my data suggests that Figure 2,
rate law 1, is the most linear, and hence the correct order of reaction can be assumed first order with
respect to Bromophenol blue. Thus, the slope of Figure 2's line is 0.00310 and therefore the k' value
for Bromophenol blue order of reaction is the slope, or 0.0031 sec–1. However, for hydroxide in
trial 1 with respect ... Show more content on Helpwriting.net ...
Also it is important to consider the spectrometers used in this lab, which had capped out above an
initial absorbance of 2 and higher in the lab, which skewed all future readings since it was
impossible to record from that value, which led to inaccurate readouts for the absorbance values. A
major source of error was that a different volumetric pipette was used to measure the initial
absorbance and in Trial 1a than the micropipette used in the other four trials. Though both
volumetric pipettes were 1 mL, the possible miniscule differences between the two glassware could
have created a significant difference in the tiny volumes. The experiment could have been improved
by use of a UV/Vis rather than a Spec, which could improve the precision and accuracy of the
measured absorbances and thus decrease the uncertainty. Additionally, measuring the absorbances
for a longer duration of time so that a significant deviation from linearity could be observed could
decrease the uncertainty in the reaction order with respect to Bromophenol blue. The best alternative
could involve measuring manipulated concentrations directly against time instead of absorbance
since the values would be in our control instead of in the machine's control, which tended to show
inevitable

The Rate Law for Chemical Reaction Among Hydrogen...
The Rate Law for Chemical Reaction Among Hydrogen Peroxide, Iodide, and Acid
To determine the rate law for a chemical reaction among hydrogen peroxide, iodide and acid,
specifically by observing how changing each of the concentrations Experiment 3 Chemical Kinetics
Objectives
1. To determine the rate law for a chemical reaction among hydrogen peroxide, iodide and acid,
specifically by observing how changing each of the concentrations of H2O2, and H+ affects the rate
of reaction.
2. To observe the effects of temperature and catalyst on the rate of reaction.
Introduction
Generally, two important questions may be asked about a chemical reaction: (1)How far do the
reactants interact to yield products, and (2) ... Show more content on Helpwriting.net ...
7. 10 cm3 of 2.0 M sulphuric acid was measured with a 10 cm3 clean measuring cylinder. It was
pour into the conical flask.
8. 1 g of solid KI (record the exact mass) and 3 drops of ammonium molybdate catalyst were added
into the conical flask.
9. The solution mixture was stirred until the KI dissolves.
10. The reaction mixture was titrated in the conical flask with the sodium thiosulphate solution until
it just turns pale yellow.
11. 3 drops of freshly prepared starch solution were added to the conical flask.
12. The titration was continued until it just changes from dark blue to colorless. The final reading
was recorded in Table 1. It was a first trial titration to estimate the volume of the sodium
thiosulphate solution required. The volume of the sodium thiosulphate solution added in titration
was calculated.

13. The given sodium thiosulphate solution was added to the burette through a filter funnel if the
volume remained was not enough to carry out another titration.
14. Steps 6–13 were repeated to obtain 2 sets of consistent results.
However, sodium thiosulphate solution was stopped draining at about 3 cm3 less than the estimated
value. Then the sodium thiosulphate solution was added drop by drop until the reaction mixture in
conical flask just changes from dark blue to colorless.
Part II) Reaction Rate Measurements
Six reaction mixtures will provide the information necessary to
determine

Taking a Look at Nucleophilic Reactions
Nucleophilic reactions occur when there is an electron pair donor and an electron pair acceptor (2).
There are two types of ways that nucleophilic reactions occur. There is the SN1 reaction and the
SN2 reaction. An SN1 is a two–step reaction that occurs when a molecule first forms a carbocation.
Once the carbocation is formed, the nucleophile comes in and attaches to the molecule (2). Below is
a general reaction scheme of an SN1 reaction:
Below is the mechanistic scheme of SN1:
In an SN2 reaction, it is a one–step reaction and occurs when a nucleophile attacks a molecule and
forces the leaving of a leaving group. Below is the mechanistic scheme of SN2: Reaction kinetics is
the study of the rates of chemical reactions. From these chemical mechanisms, one derives the rate
laws which will show how fast or slow a reaction is occuring and figure out if a first order or a
second order reaction is occuring (1). The first order reaction is an SN1 reaction. A first order
reaction has a rate proportional to the concentration of one reaction. A first order reaction formula
will be :
Rate =k[A] or rate=k[B] A second order reaction is an SN2 reaction. A second order reaction has a
rate proportional to the concentration of both reactants. The formula is as follows:
Rate=k[A][B]
SN1 reactions proceed through a carbocation, the product is a racemic mixture of the substitution
product which is a 50/50 racemic mixture. SN2 reactions do not form carbocation, but require the

The Virtue Of Order By The Reaction From The Crowd
I was disappointed not only by the difficulty of adhering to the moral virtues, but also by the
reaction from the crowd. I have now experienced the world differently, and should hope to be more
fit to follow the virtues. I have grown older, and the the world has too, so it is logical that the virtues
I mentioned previously should need to be modified.
Many virtues such as Sincerity and Cleanliness, two moral virtues which I had no trouble
accomplishing the previous time, remain applicable today. However, the virtue of Order was a bit
troublesome, so I have decided to morph it into the virtue of Planning, for one cannot, however
frustratingly, obtain complete control over the will of another. Planning in combination with
Resolution seem to account for Order. Temperance and Moderation have been combined into just
Moderation for they were awfully similar in my first experiment. I have decided to keep the other
virtues, with new meanings which ought to do a better job at describing each one then previously,
and a more lengthy explanation as to why the definition has been changed in the way it has.
The revised virtues, the new precept, and my explanation of the change are now:
Moderation Compromise by finding the middle ground; avoid the extreme. This is a mixture of
Temperance and the original Moderation virtues. I feel satisfied with their original definitions, but I
think they were too similar. This new definition takes the better half of each definition to try to be

Rate Equation And Order Of Reactant
Rate equation and order of reactant:
The rate of reaction is defined as how fast the reactant is converted to the product and it is measured
in moldm–3s–1
Rate=(change in concentration(〖moldm〗^(–3)))/(time taken (s)) , the unit of rate = moldm–3s–1
Rate = k [A]x [B]y [C]y
Orders of the reaction must be taken in to account when writing a rate equation. Where A, B and C
are reactant, x, y and z are the orders of the reaction with respect to A, B and C.
In my investigation I will vary the concentration of each reactant and I will measure time taken for
the reaction to turn colourless. I will use rate=1/time formula to calculate the rate of the reaction.
This will enable me to draw rate against concentration graph which I will use to ... Show more
content on Helpwriting.net ...
Collision theory:
Reactions only occur when particle of the reactant collide with a certain minimum kinetic energy.
This energy is also known as activation enthalpy this energy must be supplied to the reactant to
enable the bonds to stretch and break and the new bonds to form in the product. Most collisions do
not result in a reaction, only the collision with enough activation enthalpy will react to produce the
product. (1)(2) (10)
Concentration: The collision theory states that for any reaction to take place the two reactant
molecule must collide first. The rate of a reaction depends on the frequency of the collision between
the particles. Increasing the concentration of reactants increases the number of particles available to
react in the same volume of solution. More successful frequent collisions occur in a given time, with
sufficient activation energy to break the chemical bond. Therefore the rate of reaction will increase.
(1)(2)(6)
Surface area: Increasing the surface area of a solid reactant means more particles are exposed to the
other reactant, the number of reaction sites increase therefore it will react more quickly increasing
the reaction time.
Catalyst: A catalyst will increase the rate of reaction by lowering the activation enthalpy without
being chemically unchanged. They do this by providing an alternative pathway, forming an
intermediate compound with the reactant and then it forms the product from

Lab Repor Chemical Kinetics Essay
Abstract The "Chemical Kinetics" experiment was done to investigate the changes in the rate of
reaction under the effect of concentration, temperature, and presence of a catalyst. It was determined
that as the concentration of reactants and the temperature increases, the rate of the reaction increases
as well. Also, the reaction was run by the presence of catalyst, and the rate of the reaction increased
drastically in the presence of it. The order of the reaction with respect to each reactant was
calculated to be: x = 1 [I–], y = 1 [BrO3–], z = 2 [H+] by the method of initial rates. The average
rate constant was determined to be 26.7 M–3s–1, and the activation energy was calculated to be 49.6
kJ/mol. Introduction The whole purpose ... Show more content on Helpwriting.net ...
The reaction occurred under three different temperatures: 40ْ C, 10ْ C, and 0ْ C. The experiment in
this part was carried out just like part one. In flask 1, 10 mL of 0.010 M KI, 10 mL of 0.0010 M
Na2S2O3, and 10 mL of H2O were mixed. In flask 2, 10 mL of 0.040 M KBrO3 and 10 mL of 0.10
M HCl were mixed, and 3 drops of starch was added. Then, the two flasks were put in an ice bath
and cooled to about 10ْ C. Afterwards, the two solutions were mixed together while in the ice bath
with swirling until turning blue. The time was recorded. The reaction was done the same way for the
other two degrees (0ْ C, 10ْ C). If the water needed to be cold, ice was added, and if hot water was
needed, the water was heated. In the third part of the experiment, the reaction was carried out in the
presence of a catalyst. The influence of the catalyst on the rate of the reaction was investigated.
Again, mixture 1 from part one was used. In flask 1, 10 mL of 0.010 M KI, 10 mL of 0.0010 M
Na2S2O3 ,and 10 mL of H2O were mixed. In flask 2, 10 mL of 0.040 M KBrO3 and 10 mL of 0.10
M HCl were mixed, and 3 drops of starch were added. Also, one drop of 0.5 M (NH4)2MoO4,
ammonium molybdate, was added to Flask 2 which acted as a catalyst. The two solutions were
mixed until a blue color was formed and the stopwatch was stopped. All the

Blue-1 Reaction Lab Report
The Blue Dye #1 reaction can be determined by utilizing Beer's Law, and First and Second order
integration. Beer's law is first used to calculate the molar absorptivity of the Blue #1 dye. Beer's law
states that the concentration of the Blue #1 Dye is directly proportional to the absorbance of the
solution. The molar absorptivity, e was calculated to be 9.63x104 and will be used when calculating
the activation energy of the reaction. First, the rate order for both the Blue #1 and Bleach needed to
be determined. A known concentration of both Blue #1 dye and bleach are needed to conclude the
rate order of the solutions because the rate order is dependent on the concentration of the solution
used. The absorbance versus time is plotted and from ... Show more content on Helpwriting.net ...
The initial rate law is used instead. Initial Rate Law utilizes varying the concentration of the bleach
and measuring the differing initial rates of reaction and using the rate relationship shown in
Equation 5. The variable x represents the reaction order. This concludes that the reaction order of the
Commercial Bleach used in this reaction is first–order because the value of x found was 0.913. The
percent error of this first order reaction value is 9.53%. This percent error can be accounted for The
fact that both the Blue #1 dye and the bleach are first order reactions on their own, it can be
concluded that the overall reaction of the Blue #1 food dye and bleach is second order because the
exponents of each individual reaction can be added together to get the overall reaction order of the

Rate Laws Lab
Reaction order and rate laws are key to understanding the speed in which a reaction occurs and the
necessary amounts of each reactant in a reaction. Reaction order determines the concentration of
each reactant and can be used to calculate the amount of a substance in a reaction. The zeroth, first,
and second orders are the most common and were used in this lab. The order of a reaction can be
found by comparing the quantity of a specific substance and the rate in which the reaction occurred.
Rate laws contribute to the speed of differing reactions. This is a necessary principle in many fields.
For example, it is necessary to know the speed of a reaction that goes on during the inflation of an
airbag. Without knowing the rate law the speed could ... Show more content on Helpwriting.net ...
HCl concluded to be zeroth order but generally reacted faster than Na2S2O3, which was first order.
The rate law for this reaction is rate=k[Na2S2O3]. Both reactions, when graphed, showed a straight
line though. The average reaction time with HCl was 44 sec (well #1), 53.1 sec (well #2), and 61.2
sec (well #3). Resulting in a reaction rate of 0.012 sec–1 (well #1), 0.019 sec–1 (well #2), and 0.16
sec–1 (well #3). While the average reaction time for Na2S2O3 was 33.4 sec (well #4), 81.48 sec
(well #5), and 121.22 sec (well #6). Causing the reaction rate to be 0.03 sec–1 (well #4), 0.012 sec–
1 (well #5), and 0.0082 sec–1 (well #6). These results are accurate and can be considered

Reaction Rate Of Reaction
Background: A clock reaction generally involves a mixture of solutions that, after a certain amount
of time, displays a sudden colour change. This process demonstrates chemical kinetics in action,
which is the study of chemical processes and rates of reaction where the reaction rate is the speed at
which the chemical reaction proceeds. It is dependent on several factors that rely on one basic
underlying principle called collision theory. In order for a reaction to occur, the reactant molecules
must collide with each other with a certain minimum energy called the activation energy to break
and form the appropriate bonds as well as have the correct orientation when colliding. If the
favourable amount of collisions increase, then the rate of the reaction would increase. However, if
the reactant particles do not collide frequently or collide with less energy than the activation energy,
they bounce apart and the reaction would then proceed slowly or not at all. Concentration is one of
the factors that affects reaction rate. Raising the concentration of the reactants, alters the number of
particles per unit volume. Thus the more molecules present in a given volume, the greater the
probability of them colliding. If they have energies equal to or greater than the activation energy, a
higher concentration would therefore result in an increase in the reaction rate. The opposite is true if
there is a lower concentration of reactant molecules.
The relationship between rate and

Crystal Violet Lab
The purpose of this lab experiment was to investigate how fast it would take crystal violet to
decolorize. The concept of kinetics was applied by using rate to figure out the answer to the guiding
question. Rate laws display the mathematical expression of the rate of a chemical reaction and the
concentration of its reactants. Rate laws are expressed in regards to the three types of reaction
orders: 0, 1st, and 2nd. In rate laws, reaction orders are expressed as exponents and showcase the
effect of reactants' concentrations on the rate of a chemical reaction. The rate constant is displayed
as k and represents a certain reaction. Furthermore, the three rate laws are as written, rate= k for
zero order, rate= k[A] for first order, and rate= k[A]2 ... Show more content on Helpwriting.net ...
The actual rate law was rate= k[CV+]1[OH–] 1 with an overall order of 2. This was concluded by
plotting all three reaction order types in separate graphs and determining the straightest graph; then
calculating concentrations in order to determine the rate constants, pseudo k, and rate laws. For
example, M1V1=M2V2 led to [OH–] by calculating (10 ml) (0.10 M) = (20 ml) to equal 0.05 M of
OH. The actual rate constant was found by taking the negative slope of the first order graph (0.0055
s–1) over the concentration of OH (0.05 M), which equated to 0.11 M–1s–1. With this being said,
CV decolorizes at a first order rate law whose graph yields the straightest line containing the highest
R2 value (0.9973) closest to 1. Lastly, comparing the findings of the data from figures 1–4 to other
groups' date, resulted in discrepancies. Other groups claimed that CV would decolorize at a reaction
order of zero because it's graph yielded the highest R2 value. In relation to this, other groups'
findings could have been a result of certain error such as: not having the colorimeter calibrated once
the reaction mixture was prepared, using inaccurate equations to graph data, and not using graduated
cylinders to obtain precise

How Does Temperature Affect The Rate Of Reaction Experiment
From the results that were collect throughout the experimental investigation has proved the
hypothesis to only be partially right. Multiple tests were made when conducting the experiment, two
clear solutions were combined at various temperatures and concentrations. The hypothesis states that
by adjusting the concentration of the reactants will cause the reaction to either speed up at a higher
concentration or slow down at a lower concentration. In the reaction temperature should have a
similar effect on the experiment, in that increasing the temperature will cause an increase of particle
movement and cause more collision, thus increasing the reaction rate. Therefore decreasing the
temperature will decrease the rate of the reaction. From the results given in Tables 2 and 3 it shown
that every time the concentration is halved the time is increased. When the concentrations of both
KIO3 and NaHSO3 are decreased the time has increased, some concentrations having a higher
increase than others. In each concentration decrease the time is at least doubled from the previous
concentration time, which is therefore increasing the rate of the reaction. Although unlike the change
in concentration, when the temperature of the reactants were either increased or decreased their was
no significant change. As shown in table 4. When the temperature was either increased or decreased
to 10°C, 30°C, 40°C and 50°C the average times are all within one second of each other. Except for
when the reactants were cooled to a further 5°C the reaction time was increased by more than
double the reaction time. This showing that the results from the investigation did not fully support
the stated hypothesis. The standard rate law for the Iodine Clock reaction, is a first order reaction for
KIO3 and a second order reaction for NaHSO3. To determine this you must substitute the values
into the equations and whichever produces a straight line graph determines whether it is a first or
second order reaction. In graph 1. It shows that KIO3 is a second order reaction, this is due to the
second order reaction graph being straighter than the first order graph. This meaning the reaction
proceeds at a rate proportional to the square of the concentration,

Kinetics Of Acid Catalysed Propanone / Iodine Reaction
Alexander Unsworth – Tomlinson
Candidate Number: 9133
Kinetics of Acid–Catalysed Propanone/Iodine Reaction
Equation for the reaction :
CH3COCH3(aq) + I2(aq) ––> CH3COCH2I(aq) + H+(aq) + I–(aq)
Iodine + Propanone –> Iodopropanone + Hydrogen (cation) + Iodine (anion)
Introduction: Aims:
To vary the concentrations of each reactant along with the sulphuric acid in order to observe and
measure its effect on the overall rate of reaction in absorbance using colourimetry.
2) Calculate the a mean rate constant using orders of reactions and the rate equation allowing for the
overall order or reaction to be found.
3) Calculate the activation energy of the reaction using different versions of the Arrhenius equation.
4) Try and propose a mechanism for the reaction using the orders of reaction taking into account the
iodine, propanone and sulphuric acid.
Chemical Background:
Iodopropanone is formed from a redox reaction between iodine and propanone which is irreversible.
Hydrogen ions (H+) are used as a catalyst for the reactants and are disassociated from Sulphuric
acid. During the reaction the iodine solution turns from a dark brown colour into a colourless
solution this is because the iodine is reacting to produce the iodopropanone and the the colourless
iodine ions. The way I will be observing the progress of the reaction is by using a digital
colourimeter to measure the absorbance value of the change in colour.
Chemical Theory:

The Collision

The Concentration Of Reactants During The Fastest Rate Of...
It was hypothesised that the solution at 80 degrees C and the greatest concentration of reactants
would result in the fastest rates of reaction and reaction time. This was because of the scientific
reasoning that at larger temperatures and concentrations atomic particles collide with greater energy
and more frequently. The data collected from the experiment supports the hypothesis. It was
observed that the largest concentration of reactants, 0.2M and 0.5M, had the highest reaction rate of
0.00484 mol/L/s, and quickest induction phase, 1 second. Also, the slowest reaction time was 21.51
seconds and was the reaction between KIO3 at 0.05M and Na2S2O5 at 0.125M, which were the
smallest amounts of reactant. In addition, the highest temperature tested, which was 80 degrees,
produced the fastest reaction time with 4.58 seconds, while 70 degrees ranked close behind in
second. These results align with previous data and theories, which state that as temperature and
concentration rise, the time it takes to react generally decreases and the reaction rate increases
(Wilbraham, 2002), (UC Davis, n.d.).
From the testing and results several major trends and relationships were discovered. The first
observation, which supports the hypothesis, is the relationship between concentration of reactants
and reaction rate. As concentration increased, the reaction rate also increased in direct proportion.
This was shown by the results because when the concentrations doubled, from 0.1M and 0.125M to

Determination of a Rate Equation Essay
Determination of a Rate Equation
Rate equation has the form rate = k [A]x[B]y which shows how the rate of a chemical reaction
depends on the concentration of the reactants
(A&B) and the rate constant k. The rate equation normally indicates what species are involved in the
rate–determining step and how many species are involved.
A rate equation is used to describe how the concentration of a product increases or the concentration
of the reactants decreases with time, the equation also indicates how the concentration of one or
more reactants directly affects the rate. Occasionally it can even be the concentration of a product
that affects the rate. In general the rate equation for the reaction: A + B C + D
Is found by experiment to ... Show more content on Helpwriting.net ...
The equation is:
2HCl(aq) = Na2S2O3(aq) 2NaCl(aq) + S(s) + SO2 (g) = H2O(l)
APPARATUS
=========
· Conical flask.
· Beaker.
· Pipette.
· Burette.
· Clamp stand.
· Grippers.
· Funnel.
· Stop watch.
· Labels.
· Marker pen.
· White paper.
· Goggles.
SOLUTIONS
=========

· 0.2 mol dm–3 of sodium thiosulphate.
· 2.0 mol dm–3 of hydrochloric acid.
· Distilled water.
DILUTION TABLES
===============
2.0 mol dm–3 of hydrochloric acid.
––––––––––––––––––––––––––––––––––––––––––––––––––––––––
Concentration of HCl Volume of HCl Volume of water (Mol dm–3) (cm3) (cm3)
2.0 10.0 0.0 1.0 5.0 5.0 0.5 2.5 7.5 0.25 1.75 8.25 0.0 0.0 10.0
––––––––––––––––––––––––––––––––––––––––––––––––––––––––
0.2 mol dm–3 of sodium thiosulphate.
Concentration of Na2S2O3 Volume of Na2S2O3 Volume of water (Mol dm–3) (cm3) (cm3)
0.2 50.0 0.0 0.1 25.0 25.0 0.05 12.5 37.5 0.025 6.25 43.75 0.0 0 50.0
METHOD 1
========
· Firstly set

Hydrochloric Acid And Acetone Lab Report
Abstract
This experiment was conducted to find the rate law for the reaction of iodine with acetone. This was
found by using the method rates. The orders of acetone and hydrochloric acid is one while the order
of iodine is zero. The procedure was meant to notice the disappearance of one reactant, Iodine.
Introduction
A chemical reaction is when chemical substances are changed into other substances. When a
chemical reaction takes place, chemical bonds break and new ones are formed. Kinetics is the study
of the rate and mechanism of chemical reactions. Reaction mechanism is a series of individual
chemical steps by which an overall chemical reaction occurs 1. These mechanisms are important in
deciding what is the most efficient way of causing ... Show more content on Helpwriting.net ...
The method of initial rates is a common way to find the order. For this experiment, the time it takes
for the color of iodine to disappear is measured as the initial rate of reaction. HCl was used as the
acid catalyst. An acid catalyst makes a reaction happen faster, but does not get consumed in the
process. The purpose of this experiment is to find the rate law for the reaction of iodine with acetone
by using the method rates.
Experimental
First, the following materials were gathered: Four fifty milliliter beakers, four pipets, four graduated
cylinders, about 30 mL of 4.0 M Acetone, 1.0 M HCl, and 0.005 M Iodine, a squeeze bottle filled
with distilled water, a spectrometer, and a tablet that displays the data from the spectrometer. The
acetone, hydrochloric acid, and iodine were put into three of the fifty milliliter beakers. The tablet
was connected to the spectrometer via Bluetooth. The spectrometer was calibrated for a wavelength
of 410 nm. It was calibrated by filling a cuvette with water and placing it inside the spectrometer.
The room temperature was recorded. Reaction number one started off the experiment. For reaction
one, as recorded in Table 1, 3 mL of Acetone, 3 mL of HCl, 8 mL of water, and 4 mL of Iodine were
measured out into the four graduated cylinders. Then, the acetone, hydrochloric acid, and water
were combined into the last remaining fifty milliliter beaker. At the same time, the

Investigating The Reaction Rate Of Crystal Violet
orah Albaiz
CHMY143–016
Katie Link
Lab Partner: Lydia Aman
Crystal Violet Kinetics
Purpose:
The purpose of this experiment was studying the reaction rate of crystal violet with NaOH by
observing the concentration using the MicroLAB colorimeter, monitoring how the reactant
concentration affects reaction rate constant, determining the reaction order, and to calculate the
reaction pseudo rate constants and the true value rate constant. The rate of the reaction of crystal
violet with NaOH is given by the generalized rate law, rate = k [OH–]x [CV]y where k is the rate
constant for crystal violet and CV is crystal violet, C25H30N3+. Where x and y are the reaction
orders. The equation can be rewritten as:
Eq.1 k' = k[OH–]x
Since solutions of crystal violet obey Beer's law, absorptivity can be calculated using the following
equation:
Eq.2 At = ε l[CV]t
Where A is the initial absorbance when the experiment first starts, l is the path length of the cuvette
(2.54 cm), and [CV]t is the initial concentration of crystal violet.
Procedure:
Determination of Reaction Order in Crystal Violet
MicroLAB Kinetics program was opened, then the colorimeter was calibrated to a 100%
transmission by filling a marked, clean, clear cuvette, about ¾ full of deionized water. The cuvette
was wiped with a Kimwipe from the outside before putting it in the colorimeter. The cuvette was
inserted in the chamber, then the cap was closed, and the Read Blank button was clicked to start the
calibration.

Chemical Kinetics Lab Report
Review 3: Text Chemical kinetics is the study of rates and mechanisms of chemical reactions. In our
study of chemical kinetics, experimental data identifying the initial concentrations of reactants and
the instantaneous initial rates of multiple trials is used to determine the rate law for the reaction, the
order of the reactants, the overall reaction order, and the average rate constant. By comparing the
instantaneous initial rates and the initial concentrations of the reactants for two trials, it is possible to
deduce the order of each reactant. In order to determine the order of A, the two trials must be
selected such that the concentration of A changes while the concentration of B is held constant. In
this case, trial 1 and trial 3 or trial 2 and ... Show more content on Helpwriting.net ...
k=(0.103M–2s–1+0.103M–2s–1+0.103M–2s–1+0.103M–2s–1)/4=0.103 M–2s–1 The average rate
constant is 0.103M–2s–1. Combining everything, the rate law for the reaction is Rate=(0.103M–2s–
1)[A]2[B]1. We know that the reaction is 2A+2B→C+D. Based on the orders we calculated for A
and B, we know that this reaction is not an elementary reaction because both of the coefficients of A
and B are 2, which do not match the calculated orders of A and B, which are 2 and 1 respectively.
Also, if this were an elementary reaction, we would expect 3 molecules to perfectly collide with
each other, which is highly unlikely. As a result, it is more likely that there was an intermediate and
that multiple steps were involved. Through experimental data, we can not only determine the order
of reactants, but also the rate law, the average rate constant, and the overall order of the reaction.
Using the orders calculated, we can also determine the integrated rate plot that best represents the
reactants and the type of

Activation Rate Law
Abstract. The objective of this work was to determine the activation energy and rate law for the
reaction between hydrochloric acid and magnesium shot. An analysis of the reaction's initial rate at
varying molarities of hydrochloric acid and masses of magnesium shot, along with the method of
initial rates was used to determine the rate law for the reaction; rate = k (SA of Mg)a [H+]b. It was
hypothesized that the reaction would be second order overall. After determining the rate law, the
activation energy was determined by changing the temperature at which the reaction took place.
Using the rate constants calculated at the two different temperatures, the activation energy was
calculated. It was found, experimentally through the method of ... Show more content on
Helpwriting.net ...
One of the biggest sources of error was letting some of the hydrogen gas escape during the reaction
because the cap of the pressure probe could not be placed on top of the flask quick enough. Not
being able to measure all of the change in pressure resulted in a lower rate law because the graph of
pressure versus time had a smaller slope. In future experiments, a system in which the reactants
could be mixed with the pressure probe already in place would yield much more accurate results
because the change in pressure would be higher which would cause the initial rate to be larger.
Another source of error in the experiment was keeping the temperature of the trial in the ice bath
consistent. If the temperature of the HCl changed from when it was measure to when the reaction
took place, the activation energy calculations would have had increased error. Because the HCl was
still sitting in the ice bath, if the actual temperature was colder that what we calculated, the
activation energy would have been lowered. For future experiments, the temperature of the HCl
should be recorded as soon to when the Mg shot is added as possible, or the experiment could be run
in a temperature controlled

Rate Law Lab
The objective of this experiment was to determine the rate law for a chemical reaction between
crystal violet and hydroxide. A rate law is a part of kinetics, which is the study of how fast reactions
occur and how to control the rate of a reaction (4). The rate law is be determined by measuring and
graphing the absorbance of reactants during the reaction. The reaction was first order with respect to
crystal violet (CV+) and hydroxide (OH–). Since crystal violet is in much smaller concentration
than hydroxide, the experiment captured the reaction rate and order of crystal violet while the order
of OH was calculated post–lab using the pseudo first order method (eqn 1,2,3). The rate law for
CV++OH– CVOHis Rate = 0.1644m–1s–1[CV+][OH–].
Introduction
Kinetics is the study of how fast reactions occur and how to control that rate. The rate of a reaction
is defined by the change in concentration of reactants and products over time (4). The rate law for
this reaction is Rate = k[CV+]m[OH–]n, where m is the order with respect to CV+, n is the order
with respect to OH–, and k is the rate constant. Because the concentration of OH is so much larger
than the concentration of CV+, only [CV+] will change when significantly, which is why the
experiment only focuses on the change of concentration in crystal violet while the concentration ...
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A rate law is a mathematical equation that describe the rate of a reaction. A reaction rate is the
change in concentration of reactants or products over time (3). A Vernier colorimeter was used to
measure the absorbance (or concentration) of CV+ when mixed with OH. The concentration of CV+
generally decreased over time at a similar rate when mixed with 0.1M and 0.2M hydroxide. The data
was copied into Excel then made into

Thiosulfate And Iodine Reaction Lab Report
In general, this experiment was able to accomplish and establish the rates at which persulfate ions
are able to oxidize iodine ions. The reaction between thiosulfate and iodine is known to be rapid,
unless there is a sufficient amount of thiosulfate present for the iodine to react with. When
thiosulfate is reacting with iodine, the solution will remain colourless, until all the thiosulfate is used
up. After a short while, the clear solution will turn into a deep blue when iodine binds to starch. This
process allows the rate of reaction to be determined for the persulfate reaction, when the ability to
record time is accurate.
The first part of this experiment was to determine the individual concentrations of the reactants. In
order to do ... Show more content on Helpwriting.net ...
Thus, the reaction orders remain valid. Next the value of ΔS2O8–2 was retrieved by dividing
[S2O3–2] by 2, which was concluded to be 9.1x10–4 mol/L. Using the ΔS2O8–2 and the time it
took for the reaction to complete, the rate was determined.
For trial 1, the rate of reaction was 1.61x10–5 mol L–1 s–1, k was 0.167 mol2 L2 s–1 and the ionic
strength was 0.2362 mol L–1 s–1. For this reaction no electrolytes were added to speed up the
reaction. The reaction run for 1 can be considered as the model reaction for this experiment, since
other substances were not added in. For trial 2, the rate was 7.95x10–6 mol L–1 s–1, k was 0.331
mol2 L2 s–1, and the ionic strength was 0.1327 mol L–1 s–1. In this trial, the concentration of
[S2O8–2] was reduced, while (NH4)2SO4 was added. The electrolytes from (NH4)2SO4 were able
to diffuse, and replenish the reaction's ionic strength. Thus, the k value was higher and the reaction
was quick. For the third trial, the rate of reaction was 1.83x10–6 mol L–1 s–1, k was 0.305 mol2 L2
s–1, and the ionic strength was 0.1055 mol L–1 s–1. The concentration of S2O8–2 is ¼ of trial 1,
thus the rate of the reaction also decreased, and the overall reaction took longer to finish. For the
fourth

Chemical Reaction On Chemical Kinetics
Olivia Isaacs
C127
15 November 2014
Chemical Kinetics
Objective:
This experiment runs many reactions varying the concentrations of the reactants in order to
determine the order for each component and the rate constant.
Introduction:
Chemical kinetics is the study of how fast a chemical reaction occurs and the factors that affect the
speed of reaction.1 Reaction rates are the measure of how much the concentration of reactants
change during a given reaction.1 The rate of change of the reactants, Rate = – Δ [X]/Δt, is related to
the slope of the concentration vs. time graph.1 From observing reaction rates, the overall order of
the reaction and the rate constant can be calculated by using the integrated rate laws. For a zero–
order reaction, the rate law can be written as [A]t = –kt + [A]0, where [A]t is the concentration at a
given time, k is the negative slope, t is the time, and [A]0 is the initial concentration.2 Using the
same variables, a first order reaction can be written as ln[A]t = –kt + ln[A]0 and a second order can
be written as 1/[A]t = kt + 1/[A]0.2 On a graph, these concentrations are plotted vs. time, allowing
the R2 value and equation of the line to be calculated. The R2 value is used in determining the order
of the reaction. The closer the R2 value is to 1, the more likely that the graph displays the correct
reaction order. The y=mx+b equation provides information about the slope and y–intercept, essential
when determining the order and rate constant.

Lab Report On Kinetics Of Alcohol Oxidation
Lab Instructor's Name: Zhen Qiao
Student's Name: Nhu Duong
Section # CHEM 102 – 110
Experiment #6
Date of the experiment: 01/29/2016
Title: KINETICS OF ALCOHOL OXIDATION
Drexel University Winter 2016
Introduction:
This report concerns the experiment of determining the order of reaction and the kinetic rate
constant of alcohol oxidation. This experiment relates to the knowledge of chemical kinetics, the
application of Beer's Law, and other calculations.
Chemical kinetics involves the examination of reaction rates, which are the speeds of chemical
reactions. There are chemical reactions which proceed in long periods of time as well as chemical
reactions proceeding in short periods of time. Regarding reaction rates, the reaction order and
kinetic rate constant are considered.
Take a general reaction in solution for instance: aA+bB⇌Product(s) The reaction rate is calculated
using the following function: rate=k〖[A]〗^x 〖[B]〗^y
Where [A] and [B] are the concentrations of A and B in the solution, unit: mol/L or M; x and y is the
reaction order of A and B respectively; k is the kinetic rate constant, its unit depends on the order of
the reaction.
The values of x and y are the partial order corresponding to A and B. The sum of x and y is the
overall reaction order. The values of x and y can be either negative or positive. They can also be
either integer or fractional. The reaction rate can be understood as how fast the reactants are
consumed or how fast the

Compare The Rate Law For The Reaction Between Crystal...
Rate Law for the reaction crystal violet and hydroxide ions
Objective
This experiment was to determine the rate law for the crystal violet with hydroxyl ions reaction.
Materials
Crystal violet dye, sodium hydroxide (NaOH) and spectrophotometer were used in this experiment.
Theory
The rate of reaction is a topic in chemical kinetics that explains the speed at which the chemicals
interact or it is the rate at which a product is formed by a chemical reaction. There are factors that
determine the rate of a particular reaction. These factors include; reactant concentration, the
temperature of the reaction and presence or absence of a catalyst. Most chemical reactions increase
in speed as the concentration of the reactant increases. This is because ... Show more content on
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This is because the moles of crystal violet dye react with the hydroxyl ions present in sodium
hydroxide (Wheeler and Sigmann 328). The slopes of both the first order and second order graphs of
In (concentration) against time gave the rate constant (k) of each reactant. For the first order
absorbance reaction, the concentration of sodium hydroxide was 0.01M, which has a reaction rate
constant of –0.055 (mol dm–3 s–1) while the reaction rate constant of the first order absorbance
where the concentration of sodium hydroxide, 0.020M, was –0.095 (mol dm–3 s–1). When rate
constant (NaOH 0.01M) was divided by (NaOH 0.020M) gave 0.579 (mol dm–3 s–1). The
functional change of y as the concentration rate changes is calculated to be 0.8. Therefore, the order
of reaction with respect to the hydroxyl ions was 0.8.
Conclusion
In this experiment, the reaction order with respect to hydroxyl ions was calculated. The comparison
of rate 1 (initial concentration) and rate two (second concentration) showed that the value of y which
is the rate of reaction in respect to hydroxyl ions was 0.8. Therefore, the order of reaction with
respect to hydroxide ions was 0.8. By using the graphical method, the value of x=1. x represents the
order of reaction with respect to dye(crystal violet). The overall rate law was found to be; Rate=k
[crystal violet] 1[hydroxide ions]

What Is A Pseudo-Order Kinetic Reaction
In the experiment, I found that the reaction is a pseudo – first order reaction from analyzing the
graph (Figure 2,3,4) and determine which were closer to a straight line. I found the pseudo– order
kinetic rate constant that correspond to it is 0.0016 1/s because it was the slope of the straight line
(Figure 3) Using this pseudo–order kinetic rate constant I found the real constant of the rate law by
dividing the concentration of ethanol which gave us the constant to be 2.66×〖10〗^(–3) L/(mol s) .
Overall, during the lab there should not have been too much error. This lab given us the opportunity
to revised the data if necessary before continue to the next step. When we were asked to construct a
Beer's Law Calibration Curve we created a curve

The Effect Of A Clock Reaction On The Reaction Rate Of A...
A clock reaction is a form of chemical reaction where multiple reagents are placed together with no
apparent change after a time lag until a dramatic colour change appears – called the induction period
(King, Preece and Billingham, 1999). This abrupt colour change can be due to the limiting reagent
being fully consumed which allows the remaining chemicals to react resulting in a colour change.
Specific kinetic and stoichiometric conditions are required for this standard clock behaviour to
occur. Clock reactions can be highly useful to investigate rates of reaction as the rate of the colour
change may differ due to varying factors.
The reaction rate of a chemical reaction can be defined as the speed in which the reaction occurs.
The speed ... Show more content on Helpwriting.net ...
Figure 1: Effect of order on the reactant (UCDAVIS 2, 2017)
Typically, the order of a reactant does not go beyond a second order. The greater the order is, the
greater the increase in reaction rate when the reactant is doubled. The sum of the order of reactants
for the experiment is then the order of the overall reaction (UCDAVIS 2, 2017).
The two methods to determine the rate law of a kinetics experiment is by initial rates and by
integrated rate law. To utilise the initial rates method, multiple experiments can be conducted with
differing concentrations and measuring the initial rate for the reactions (Yamauchi, Nord and
Schullery, 2016). From this, the reaction order for the corresponding substance can be calculated
and the rate constant can be found. However, the method of initial rates requires the chemical
reaction to be fairly slow and since several experiments need to be executed, it is not an ideal
method.
The integrated rate law method utilises integrated rate laws in comparison with the collected data
relating concentration and time during a kinetics experiment (SparkNotes, 2017). This method finds
the order of the reactant by plotting the time versus a function of the concentration data and using
trial and error with a selected rate law. If the integrated rate law is correct, the data points will
appear linear and if it is not linear then

Maths Level Internal Assessment : Elimination Of Drug From...
MATHS STANDARD LEVEL INTERNAL ASSESSMENT ELIMINATION OF DRUG FROM
THE BODY So, the other day I was watching a show named Grey's anatomy; it revolves around a
medical perspective, this show basically premiers how a hospital works and its base that lies on the
ground rule of saving people's life. This somehow inspired me to study science in high school so
that I could pursue some career in medical field and save someone's life. Having thought of this I sat
down wondering what I could possibly do for this assignment, and as I turned around , gazing at the
T.V I saw that the nurse was giving some kind of a drug to the patient to relive the pain. This struck
me and that's how I came to a point of utilizing mathematical procedure in pharmacokinetics (the
branch of pharmacology concerned with the movements of drugs within the body i.e. the study of
elimination and absorption of drugs by the body).I will be evaluating how the drug that is
administered to a patient is metabolised and eliminated? More concerns will lie towards the
pathways the drug chooses to lose its concentration and I will even be looking at the rate of
reactions of these drugs in the body. There are usually five steps in the process of pharmacokinetics:
Liberation: Firstly the drug that is administered to the patient is released from the formulation.
Absorption: The drug is then absorbed by the body. Distribution: The plasma helps in the
distribution of the drug throughout the body. Metabolism: The drug is

Bleaching Of Blue Lab Report
Determination of the Rate Law for the Bleaching of Blue One Objective
The objective of this investigation is to determine how the concentrations of the reagents in the
reaction between Blue One dye and Sodium Hypochlorite (bleach) will affect the rate for the
reaction. This will be done by through the determination of the rate law with respect to the orders of
reaction for the reagents as well as the overall order of the reaction.
Introduction
The question for this investigation came about when my class learned about how colorimetry can be
used to detect the kinetics of a reaction. I find this process engaging since I appreciate how
colorimeters represent the human perception of colour; likewise, my interest in how dye can be
altered in ... Show more content on Helpwriting.net ...
If the natural log of the absorbance creates a linear trend, then the solution has an order of reaction
of one; meaning that doubling the solution of the dye, in turn, doubles the rate. If the inverse of the
absorbance values (1/absorbance) creates a linear trend, then the order of reaction in respect to the
dye is two. Meaning that doubling the concentration of the dye, quadruples the rate of the reaction.
Which transformation renders the most accurate linear trend will be calculated by plotting the
variables against time and finding the coefficient of determination (R2 value) for the line of best fit.
This values states how much of the variance in the y–variable, the natural log of the absorbance or
the inverse of the absorbance, is accounted for by the least squares regression line based on the x
variable– time.
To find the order of the reaction in respect to bleach, the volume of the bleach can be doubled and
the rate of the reaction examined. This is because the concentration of bleach cannot be detected by
the colorimeter; hence, experimental rate values are to be determined. The rate is calculated by
taking the slope of the least squares regression line based on whichever y variable results in the
closest linear

The Theory And Factors Affecting Reaction Rates
Chemical Background Hydrogen peroxide is a by–product of many reactions that occur within the
body – however, it is toxic so needs to be broken down. The equation for this decomposition is as
follows: 〖2H〗_2 O_2 → 〖2H〗_2 O+ O_2 In the body, this reaction can be catalysed by the
enzyme catalase. Catalase is not removed or used up in this reaction, and speeds up the rate of
reaction. It is acting as a catalyst. The decomposition can be catalysed by other catalysts, however,
and this is the basis of my investigation. By using different catalysts, I can investigate how the rate
of reaction changes with each one and find out if catalase is the best catalyst for this decomposition.
The collision theory and factors affecting reaction rates For a reaction to take place, two particles
must collide with each other so they come in to contact. However, just colliding with each other
does not initiate a reaction. Something called the activation enthalpy must be overcome. The
activation enthalpy is the minimum (kinetic) energy required by a pair of molecules that are
colliding before a reaction can occur. So, for a reaction to take place, pairs of molecules must collide
with enough energy to equal or overcome the activation enthalpy. There are several factors that
affect the rate of reaction (three of which I am investigating). The factors are as follows:
Concentration of reactants – This factor is explored in more depth in a later section. Temperature –
this factor is explored in

Chemical Reactions And Factors That Affect The Rate Of A...
Reactions occur everywhere, and they may take decades, such as fossils, or only seconds, such as
lighting a match to occur. "Chemical kinetics concerns the rates of chemical reactions" and what
factors affects these rates (Iodine Clock, 2017). "Temperature, concentration, pressure of reacting
gases, surface area of reacting solids and the use of catalysts are all factors which affect the rate of a
reaction" (Bbc.co.uk, 2017). This is because they affect the reaction roles and yields of activation
energy, product management and reactant management (William, 2017). This is accomplished by
"making changes to the concentration, pressure or temperature of a reaction to alter the position of
the equilibrium" (Bbc.co.uk, 2017).
Accordingly, the ... Show more content on Helpwriting.net ...
It was further identified that typically for every 10oC temperature rise, the reaction rate would
double (Del Mundo et al., 2016) (Clark, 2017).
Throughout the Landolt Iodine Clock Reaction, chemical equilibria are clearly demonstrated
through the delayed colour change of combining two colourless solutions to form a dark blue
solution. The overall chemical reaction for this is as follows:
2IO–3(aq) + 5HSO–3(aq) + 2H–(aq) → I2 + 5HSO–4(aq) + H20(l)
Iodate ions + Bisulfite ions + hydronium ions → Iodine + Bisulfate + water (Iodine Clock, 2017)
The individual steps of the reaction are:
(1) IO3–(aq) + 3HSO3–(aq) → I–(aq) + 3SO4–(aq) + 3H+(aq)
Iodate + Bisulfite → Iodide + Sulfate + Hydronium ions
(2) IO3– (aq) + 5I–(aq) + 6H+aq) → 3I2(aq) + 3H2O(l)
Iodate + Iodide + Hydronium ions → Iodine + Water
(3) I2(aq) + HSO3–(aq) + H2O(l) → 2I–(aq) +SO4–(aq) +3H+(aq)
Iodine + Bisulfite + Water → Iodide + Sulfate + Hydronium ions
(4) I2(aq) + I–(aq) + starch → dark blue starch–I3– complex
Iodine + Iodide + starch → dark blue starch–I3– complex
(Iodine Clock Reaction, 2015)
The first reaction is the rate determining step, and hence, determines the order of the reaction. The
equilibria is delayed in this reaction due to any of the iodine produced, immediately reacting with
any bisulfite still present, before then being converted into colourless iodide (reaction 2). Therefore,
to demonstrate this reaction, the

Practical Lab Report On Rate Of Reaction
Practical lab book
Changing the conditions of a reaction and the effects those changes have
Abstract
The purpose of this experiment is to find out the rate of reaction, the reaction constant, rate law and
also to change the temperature and concentration of the reactants in the reaction, determine the
reaction rate and find out what effects these changes have on the reaction. The reaction is measured
by the time taken for a colour change to be observed.
The rate of reaction is affected by the change in concentration and change in temperature of the
reactants. The result of the experiment showed that the reaction order of the [S2O82–] was 1st order
and [I–] was 3rd order.
Introduction
Rate of reaction is a way of determining the speed at which a reaction proceeds. It can be calculated
by measuring the appearance of a product or the disappearance of a reactant over time. Reactions
occur when the particles of the reactants collide successfully. The more successful collisions
between particles, the faster the reaction rate. There are factors that affect collision theory and
consequently the rate of reaction. High concentration of reactants results in more particles and a
higher likelihood of successful collisions. Higher temperature of the reaction also increases the rate
of reaction because the energy of the particles is higher which again increases the chances of
successful collisions. (Lawrie RYAN, 2000). On completion of this experiment, the reaction orders
can be

Peroxydisulfate And Iodide Clock Reaction
Individual Investigation
Title
Aim
To investigate the effects of changing the concentration of the reactants and catalysts in the
peroxydisulfate/iodide clock reaction. The six investigations differ by either changing the
concentration of reactants (potassium peroxydisulfate and potassium iodide), changing the
concentration of catalyst, changing the catalyst used, and by changing the temperature of the
solutions. The clock reaction method was used as a way of measuring the rate of reaction so I could
see how the varying concentrations and temperatures affected the rate of reaction.
Introduction
I am investigating the peroxydisulfate/iodide reaction. This is a clock reaction and so involves
measuring the rate of reaction, whilst observing how ... Show more content on Helpwriting.net ...
The sum of x and y will give the overall reaction order. The orders of reaction are based on the
kinetics of the reactants and so can only be found after the experiments have been completed. The
can be worked out using the initial rate method, in which experiments are repeated using different
concentrations, then the initial rate of the reaction is calculated (using 1/time as shown above). The
initial rate is then plotted against the concentration on a graph, showing the order of the reaction.
There are four known orders of reaction; zero order, first order, second order, and third order. Zero
order reactions have a rate that is not dependent on the concentration of the reactants. Therefore, this
is easily identifiable on an initial rate against concentration graph as it is shown as a straight,
horizontal line (see Figure 1). So, if the concentration of the reactant/s is increased, the initial rate of
reaction will not change, and so the concentration of the reactant is proportional to the time (Rate =
k). If the orders of reaction with respect to each reactant used in the peroxydisulfate/iodide clock
reaction are zero order, they can be written as the

Chemical Reaction Of Chemical Reactions
Chemical kinetics is the study of rates during chemical processes and the speed at which they occur
(Chm.Davidson, 2016). Chemical kinetics can be altered by the effect of various variables and the
re–arrangement of atoms. An example of kinetic processes can be seen in many experiments such as
the 'Landolt Iodine Clock Reaction.'
Clock reactions represent chemical reactions in which two colourless solutions are mixed together;
at first no reaction takes place but after a short period of time the solution can be seen to undergo a
change in colour. Within the Landolt iodine clock reaction Potassium Iodate and Sodium Bisulphite
react to yield iodine, which in return reacts with the starch molecules to form the blue solution
(RSC, 2016). The process of this experiment can be explained through the following chemical
reactions:
The iodine ions are produced, due to following reaction between iodate and bisulphite:
IO3− + 3 HSO3− → I− + 3 HSO4−
2. The iodate that is left in excess after the first reaction will oxidise with iodide formed to generate
iodine:
IO3− + 5 I− + 6 H+ → 3 I2 + 3 H2O
3. Instantaneously the iodine is reduced back to iodide by the bisulphite I2 + HSO3− + H2O → 2 I−
+ HSO4− + 2 H+
Consequently the iodine will react with the starch to create the coloured solution; this will only
occur once the bisulphite is fully consumed (Unomaha, 2016). Throughout the course of these
chemical reactions, the rate at which the reaction occurs takes place

Essay on Computerized Data Acquisition of a Second Order...
Computerized Data Acquisition of a Second Order Reaction Abstract Introduction The rates at
which reactions occur depend on the composition and the temperature of the reaction mixture.
Usually the rate of reaction is found to be proportional to the concentrations of the reactants raised
to a power.1 There are many reactions that have a rate law in the form of: (1) v = k[A]a[B]b
According to reference1 the power to which the concentration of a species (product or reactant) is
raised in a rate law of this nature is the order of the reaction with respect to that species. In equation
(1) first order with respect to [A] and first order with respect to [B]; however, the overall reaction is
the sum of the individual orders. Thus ... Show more content on Helpwriting.net ...
The Erlenmeyer flask was swirled for 2–3 seconds before pouring the reacting mixture into a 1–cm
cuvette. The cuvette was conditioned with the reacting solution 4 times before being placed into the
sample holder of the spectrophotometer. An absorbance reading was taken at 30 seconds and every
30 seconds thereafter for a total of 6 minutes. The same process was implemented with the Cary 50
Bio except that each sample was analyzed by the computer for 7 minutes and 53 seconds.
Data/Results 0.025 M NaNO3 0.05 M NaNO3 Time (sec) Asb [K3Fe(CN)6] [C6H8O6] Asb
[K3Fe(CN)6] [C6H8O6] 30 0.622 0.0006146 0.0003573 0.653 0.00065 0.0003726 60 0.617
0.0006097 0.0003548 0.640 0.00063 0.0003662 90 0.606 0.0005988 0.0003494 0.628 0.00062
0.0003603 120 0.600 0.0005929 0.0003464 0.619 0.00061 0.0003558 150 0.593 0.0005860
0.0003430 0.609 0.00060 0.0003509 180 0.584 0.0005771 0.0003385 0.600 0.00059 0.0003464 210
0.578 0.0005711 0.0003356 0.591 0.00058 0.0003420 240 0.571 0.0005642 0.0003321 0.583
0.00058 0.0003380 270 0.564 0.0005573 0.0003287 0.575 0.00057 0.0003341 300 0.559 0.0005524
0.0003262 0.567 0.00056 0.0003301 330 0.552 0.0005455 0.0003227 0.560 0.00055 0.0003267 360
0.546 0.0005395 0.0003198 0.553 0.00055 0.0003232 0.01 M NaNO3 0.2 M NaNO3 Time (sec)
Asb [K3Fe(CN)6] [C6H8O6] Asb

Crystal Violet Formal Lab Essay
Determination of Reaction Rate Law from the Reaction of Crystal Violet with Sodium Hydroxide
______________________________________________ Abstract: This experiment helps determine
the rate of reaction of crystal violet while it reacts with sodium hydroxide with respect to crystal
violet. The amount of sodium hydroxide is varied in this experiment while crystal violet is kept at a
constant. The transmittance of crystal violet is observed and recorded using a colorimeter and the
data obtained is used to plot graphs which are manipulated using LoggerPro software to produce the
desired outcome; rate of reaction of crystal violet. Upon completion of the experiment it was seen
that the rate of reaction of crystal violet turned out to be 1 ... Show more content on Helpwriting.net
...
It must be noted that only the absolute value of the slope matters in this situation. Third order
reactions have somewhat a similar story except they require a plot of 1/concentration versus time to
determine rate of reaction. When all three graphs are plotted, the graph with the line of best fit, or
the one in which all point seem to be on a straight line is the correct one for the reaction. This is
easily drawn using the LoggerPro software. When all three graphs are drawn, the graph with the best
fit line and lowest root mean square error, or the lowest deviation from the best fit line, is the graph
to be used to determine reaction kinematics. This knowledge is acquired from the equations of the
integrated rate laws which are explained in the textbook. The solutions are mixed in small amounts
in cuvettes and inserted into the colorimeter, which reads the percentage transmittance during the
time period. The colorimeter has an enclosed space for the cuvette to be inserted making sure light
from other sources does not interfere with the reaction, hence providing accurate results. The rate of
the reaction is determined by using the equation: Rate= k [CV+] [OH–], where k is the rate constant
for the reaction. Materials: Solutions of crystal violet and sodium hydroxide were available in the
laboratory which were previously prepared of concentrations 2.00 E–5 and 2.00 E–2 respectively.
Deionized water was used

Enzyme Reaction Lab
Introduction To maintain basic life functions, cells have to perform reactions. Early chemists found
that whenever certain chemicals were present in reactions, substances were changed from one form
to another. These certain chemicals present in these reactions were named catalysts. These early
chemists also found that enzymes were not consumed or altered in the reactions. Enzymes also
lower the activation energy in cells causing the reactions to occur at a faster past. In our cells,
enzymes are proteins that act as these catalysts. Enzymes cause cells to perform basic life functions
at a speed that makes life possible. Enzyme reactions occur in nature, and they are spontaneous
reactions. (Stegenga Biology 101).
Enzyme names are based ... Show more content on Helpwriting.net ...
The wavelength was set to 540 nanometers. It was set to 540 nanometers in order to produce a green
light that would be absorbed by the benzoquinone for the reasons previously stated. There were five
different tube samples that each held a total of 5mL of liquid. 1 mL of enzyme and 2 mL of EDTA
were added to Tube 1. 1 mL of enzyme and 2 mL of PTU were added to Tube 2. 1 mL of enzyme
and 2 mL of Citric Acid were added to Tube 3. 1 mL of enzyme and 2 mL of dH2O were added to
Tube 4. Tube 4 was the control tube considering it contained no chelating agents. With this, the
normal rate of absorbance could be measure and compared to those that contained chelating agents
in order to see if the ion that was taken away by the chelating agent was needed. 5 mL of dH2O was
added to Tube 5. Tube 5 was the calibration tube considering it contained nothing but distilled water.
All of the substances were added to the tubes by pipettes in order to accurately measure the amount
of substances. Whenever the chelating agents (the PTU, Citric Acid, and EDTA) were added to the
tubes, the agents were taken from the top of the solution by a pipette in order to avoid the parts of
the solution that had settled. The tubes sat at room temperature for 10 minutes, and they were
inverted and shaken every 2 minutes in order to mix the enzymes and chelating agents. They sat at
room temperature for 10 minutes while being mixed every 2

Crystal Violet Kinetics Lab

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Crystal Violet Kinetics Lab