This lab uses a color-fading reaction to measure reaction rate. Crystal violet dye reacts with sodium hydroxide (NaOH) and slowly loses its purple color. You track that color change using absorbance measurements over time, then use those measurements as a stand-in for concentration. From there, you figure out the order of the reaction and write the rate law.

The key move here is connecting spectroscopy to kinetics. Absorbance tells you how much light a solution absorbs, and that links directly to concentration through Beer's Law. So instead of measuring concentration directly, you measure light, and the math does the rest.

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Why This Lab Matters for the AP Exam

Kinetics is one of the most heavily tested units in AP Chemistry. The free response section regularly asks you to determine reaction order from data, write rate laws, and interpret graphs of concentration vs. time. This lab gives you hands-on experience with all of that.

It also connects spectroscopy (Topic 3.11) to kinetics (Unit 5) in a way that shows up on the exam. You need to understand why visible light works for this particular reaction, and what absorbance actually tells you about what is happening at the molecular level.

CED Connections

Topic 3.11: Spectroscopy and the Electromagnetic Spectrum

Learning Objective 3.11.A asks you to explain the relationship between a region of the electromagnetic spectrum and the types of molecular or electronic transitions associated with that region.

In this lab, you use a colorimeter or spectrophotometer that works in the ultraviolet/visible region of the electromagnetic spectrum. Crystal violet absorbs visible light because it undergoes electronic transitions (electrons jumping between energy levels). When the dye fades, those electronic transitions stop happening, and absorbance drops. That is essential knowledge 3.11.A.1c in action.

The other two regions matter for context:

Microwave radiation is associated with transitions in molecular rotational levels (3.11.A.1a)
Infrared radiation is associated with transitions in molecular vibrational levels (3.11.A.1b)

You are not using those regions in this lab, but you should be able to explain why visible light is the right tool here and not IR or microwave.

Topic 4.3: Representations of Reactions

Learning Objective 4.3.A asks you to represent a chemical reaction with a consistent particulate model.

The reaction between crystal violet and NaOH has a balanced equation with specific coefficients that tell you the ratio of particles reacting. You should be able to draw a particulate drawing that shows the reactant molecules colliding and the product forming, with particle counts that match the stoichiometry. This is essential knowledge 4.3.A.1.

Topics 5.4 and 5.5: Elementary Reactions and the Collision Model

Learning Objective 5.4.A connects to how you write a rate law from the reaction's stoichiometry (for elementary steps). Learning Objective 5.5.A connects to why reactions happen at the rate they do.

The crystal violet reaction lets you test whether the reaction behaves as a first-order process with respect to the dye. You are also working with ideas from collision theory: molecules need to collide with enough energy (the activation energy) and the right orientation to react. The rate you measure reflects how often those successful collisions happen.

What You Need to Be Able to Do

Use Beer's Law to connect absorbance measurements to relative concentration
Plot absorbance (or ln[absorbance]) vs. time and interpret the shape of the graph
Determine the order of the reaction with respect to crystal violet by analyzing which graph gives a straight line
Write the rate law expression for the reaction
Calculate the rate constant from the slope of the linearized graph
Explain why NaOH is treated as a constant (pseudo-first-order conditions)
Connect the fading of color to electronic transitions in the visible region of the electromagnetic spectrum
Draw a particulate drawing that represents the reaction at the molecular level

Core Concepts

Beer's Law and Absorbance

Beer's Law (also called the Beer-Lambert Law) states that the absorbance of a solution is directly proportional to the concentration of the absorbing species. Written out:

$A = \varepsilon l c$

Where:

$A$ = absorbance (no units, it is a ratio of light intensities)
$\varepsilon$ = molar absorptivity (a constant for a given substance at a given wavelength)
$l$ = path length of light through the solution (usually 1 cm)
$c$ = concentration of the solution

Since $\varepsilon$ and $l$ stay constant during the experiment, absorbance is directly proportional to concentration. That means you can use absorbance as a proxy for concentration without ever calculating the actual molarity. When absorbance drops by half, concentration has dropped by half.

Why Crystal Violet Absorbs Visible Light

Crystal violet is a large organic dye molecule with a system of alternating single and double bonds (called a conjugated system). The electrons in that system can absorb photons of visible light and jump to higher electronic energy levels. This is a UV/visible electronic transition, which is why you use a colorimeter set to visible wavelengths (around 590 nm for crystal violet).

When crystal violet reacts with NaOH, the structure of the molecule changes and it can no longer absorb visible light the same way. The solution fades from purple to colorless. The rate of that fading is what you are measuring.

This connects directly to the electromagnetic spectrum. Infrared radiation would interact with molecular vibrational levels (bond stretching and bending). Microwave radiation would interact with molecular rotational levels. But electronic transitions require the higher energy of UV/visible photons, described by the equation:

$E = h\nu$

Where $E$ is the energy of the photon, $h$ is Planck's constant, and $\nu$ (nu) is the frequency of the radiation. Higher frequency means higher energy, which is why visible light can drive electronic transitions while microwaves cannot.

Integrated Rate Laws

The integrated rate law connects concentration to time. Which form you use depends on the order of the reaction.

For a first-order reaction:

$\ln[A] = -kt + \ln[A]_0$

This is linear. If you plot $\ln[A]$ vs. time and get a straight line, the reaction is first order. The slope equals $-k$ .

For a second-order reaction:

$\frac{1}{[A]} = kt + \frac{1}{[A]_0}$

This is also linear, but only if you plot $\frac{1}{[A]}$ vs. time.

For a zero-order reaction:

$[A] = -kt + [A]_0$

A plot of $[A]$ vs. time is linear.

In this lab, you substitute absorbance for $[A]$ since they are proportional. So you will test plots of $A$ vs. time, $\ln(A)$ vs. time, and $\frac{1}{A}$ vs. time to see which one is linear.

Pseudo-First-Order Conditions

The actual rate law for this reaction involves both crystal violet and NaOH. But NaOH is used in a huge excess compared to crystal violet. That means the concentration of NaOH barely changes during the reaction. You treat it as a constant, which simplifies the rate law so it only depends on the concentration of crystal violet. This is called a pseudo-first-order condition.

Collision Theory and Activation Energy

Collision theory says that for a reaction to happen, particles must collide with enough energy to overcome the activation energy barrier and with the correct orientation. Most collisions do not lead to a reaction because the particles do not have enough energy or are not lined up correctly.

The Maxwell-Boltzmann distribution shows how particle energies are spread out in a sample. Only the fraction of particles with energy greater than or equal to the activation energy can react. At higher temperatures, that fraction increases, which is why reactions go faster when you heat them up.

Molecularity refers to the number of particles involved in a single elementary step. A bimolecular step involves two particles colliding. A termolecular step (three particles colliding at once) is rare because the probability of three particles meeting simultaneously is very low.

Half-Life

The half-life of a first-order reaction is the time it takes for the concentration (or absorbance) to drop to half its starting value. For a first-order reaction, the half-life is constant and does not depend on starting concentration:

$t_{1/2} = \frac{0.693}{k}$

You can check your rate constant by calculating the half-life and seeing if the absorbance data actually cuts in half at that time interval.

How the Lab Works

You are watching a reaction happen in real time by tracking color. Crystal violet starts as a deep purple solution. When you mix it with NaOH, the reaction begins and the color slowly fades. A colorimeter or spectrophotometer measures how much light the solution absorbs at regular time intervals.

The logic is straightforward. More dye in solution means more light absorbed. As the dye reacts and disappears, absorbance drops. Since absorbance is proportional to concentration (Beer's Law), you are effectively tracking concentration over time without a titration or any other direct measurement.

NaOH is present in large excess, so its concentration stays essentially constant throughout. That means the only variable driving the change in rate is the crystal violet concentration. This is the pseudo-first-order setup.

You collect absorbance readings every few seconds or minutes (depending on your setup) until the reaction is mostly complete. Then you use those data points to figure out the order of the reaction and the rate constant.

The reaction itself involves crystal violet (a large organic cation) reacting with hydroxide ions. The OH- attacks the central carbon of the crystal violet molecule, disrupting the conjugated electron system that gives the dye its color. The product is colorless. This is a real chemical change, not just a physical one, and you should be able to represent it with a balanced equation and a particulate drawing showing the reactant particles and product particles.

Data and Analysis Moves

Setting Up Your Data Table

Your raw data will be time (in seconds) and absorbance. From those two columns, you calculate:

$\ln(A)$ for each time point
$\frac{1}{A}$ for each time point

Graphing to Determine Reaction Order

Plot all three graphs:

$A$ vs. time (tests zero order)
$\ln(A)$ vs. time (tests first order)
$\frac{1}{A}$ vs. time (tests second order)

The graph that gives the best straight line tells you the order. For most crystal violet experiments, the $\ln(A)$ vs. time graph is linear, confirming first-order behavior with respect to crystal violet.

Finding the Rate Constant

Once you identify the linear graph, find the slope using a best-fit line. For a first-order reaction:

$\text{slope} = -k$

So $k = -\text{slope}$ . The units of $k$ for a first-order reaction are $\text{s}^{-1}$ .

Writing the Rate Law

With pseudo-first-order conditions and NaOH in excess, your rate law looks like:

$\text{rate} = k'[\text{crystal violet}]$

Where $k'$ is the pseudo-first-order rate constant you calculated from the slope. If you want to write the full rate law including NaOH, you would need to run the experiment at multiple NaOH concentrations to find the order with respect to OH-.

Controls and Variables

Independent variable: time
Dependent variable: absorbance (which represents concentration)
Controlled variables: temperature, NaOH concentration, path length of the colorimeter, wavelength of light used

If your school runs the experiment at multiple temperatures, you can compare rate constants and use the Arrhenius equation to estimate activation energy. That is an extension, but it connects directly to Topics 5.4 and 5.5.

Checking Your Work with Half-Life

Use your rate constant to calculate the expected half-life. Then go back to your data and check whether absorbance actually drops to half its initial value at that time. If it does, your rate constant is consistent with the data.

Particulate Drawings

For Topic 4.3, you should be able to draw before-and-after particle diagrams. Show crystal violet molecules and hydroxide ions as reactants, and the colorless product as the result. Particle counts in your drawing should match the stoichiometric coefficients in the balanced equation.

Common Mistakes

Confusing absorbance with transmittance. Absorbance and transmittance move in opposite directions. When the solution gets lighter, transmittance goes up and absorbance goes down. Always use absorbance for Beer's Law calculations, not transmittance.

Forgetting that absorbance is a proxy for concentration. You are not measuring concentration directly. You are using Beer's Law to justify treating absorbance as proportional to concentration. If an exam question asks you to explain this step, you need to cite Beer's Law explicitly.

Misidentifying the order from a curved graph. If your $A$ vs. time graph curves downward, that does not automatically mean it is first order. You have to check which linearized graph is actually straight. Do not guess based on the shape of the raw data.

Treating the slope as positive. The slope of $\ln(A)$ vs. time is negative because concentration is decreasing. The rate constant $k$ is always positive, so $k = -\text{slope}$ .

Mixing up molecularity and reaction order. Molecularity describes the number of particles in a single elementary step. Order of the reaction is determined experimentally from rate data. They are only equal when you are dealing with an elementary reaction. For an overall reaction with multiple steps, you cannot assume they match.

Forgetting why NaOH is treated as constant. The pseudo-first-order simplification only works because NaOH is in large excess. If the concentrations were comparable, you could not simplify the rate law this way.

Confusing the electromagnetic spectrum regions. Visible light drives electronic transitions. Infrared drives vibrational transitions. Microwave drives rotational transitions. Do not mix these up on the exam. The question will often give you a wavelength or region and ask what type of transition it corresponds to.

Drawing incorrect particulate diagrams. Your particle counts must match the balanced equation. If the equation shows a 1:1 ratio of crystal violet to OH-, your drawing cannot show two OH- ions reacting with one crystal violet molecule.

Quick Review Checklist

Beer's Law ( $A = \varepsilon l c$ ) lets you use absorbance as a stand-in for concentration because the two are directly proportional
Crystal violet absorbs visible light due to electronic transitions; this is why you use a colorimeter set to visible wavelengths, not IR or microwave
To find reaction order, plot $A$ vs. time, $\ln(A)$ vs. time, and $\frac{1}{A}$ vs. time; the linear graph identifies the order
For a first-order reaction, the slope of $\ln(A)$ vs. time equals $-k$
NaOH is in large excess, creating pseudo-first-order conditions so the rate law simplifies to depend only on crystal violet concentration
The half-life of a first-order reaction is constant: $t_{1/2} = 0.693/k$
Collision theory explains why only a fraction of collisions lead to products: particles need sufficient energy (at least the activation energy) and correct orientation
Particulate drawings must be consistent with the balanced equation, with particle counts matching stoichiometric coefficients