Forensic Applications of 1412 Equilibrium Calculations 2026 L26
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Overview
Acid-base equilibrium and Le Châtelier's principle provide practical tools for preparing forensic samples, predicting whether drug molecules are ionized, and separating compounds before instrumental analysis. Worked examples use amphetamine, naproxen, codeine, quinine, and heroin to connect pH and pKa calculations with drug absorption models and pH-controlled liquid-liquid extraction.
Key takeaways
- A forensic sampling plan for 100 seized items can select 20 at random and then 10% of the remaining 80, producing a 28-item sample rather than analyzing every item.
- The approximation x is negligible is not justified in the 0.025 M amphetamine example because Cbase/Kb is about 39, below the lecture's threshold of 100; the quadratic formula gives roughly 14% ionization.
- For weak acids, the ionized fraction is 1 / (1 + 10^(pKa − pH)); pH equal to pKa gives 50% ionization, and higher pH favors the charged A− form.
- At pH 2, naproxen with pKa 4.2 is mostly neutral and favors the organic phase, while codeine with pKa 8.2 is protonated and remains in the aqueous phase.
- A pH near 9 can separate heroin from quinine in the worked example: heroin with pKa 7.9 is mostly neutral, while quinine with pKa values 5.1 and 9.7 remains sufficiently protonated to favor the aqueous layer.
Chapters
- Field samples may contain solvents, biological material, or mixtures of drug and adulterants, so filtration, wipes, solvent extraction, and homogenization can prepare material for analysis.
- A cited sampling plan tests all items when there are 20 or fewer; for larger groups, it tests 20 randomly selected items plus 10% of the remainder.
- For 100 seized bags, the example plan samples 20 bags and then 8 of the 80 remaining, for 28 total.
- A seized powder can contain a controlled drug, caffeine, and inert fillers such as coffee creamer, so the analyte must be separated from the sample matrix.
- Equilibrium chemistry and Le Châtelier's principle can shift compounds between phases, helping separate analytes from interfering substances.
- The lesson frames pH control as a way to prepare complex samples for sensitive laboratory instruments.
- Amphetamine is treated as a weak base: its reaction with water forms protonated BH+ and OH−, and its pKa is converted to the corresponding Kb for the ICE-table calculation.
- In the 2 M example, the calculated ionized fraction is about 0.5%, illustrating that a small amount of ionization can still affect solution pH.
- At a lower concentration of 0.025 M, the ratio Cbase/Kb is about 39, below the lecture's 100 threshold for safely assuming x is negligible.
- The lower-concentration example therefore uses the quadratic formula and gives about 14% ionization.
- At pH 2, hydronium concentration is 0.01 M; the example shifts amphetamine toward its charged BH+ form, with the calculation yielding essentially complete protonation.
- Charged forms are generally more water-soluble than their neutral counterparts, linking ionization equilibria to movement between aqueous body fluids and other environments.
- The lesson uses stomach and intestinal pH as a model for how acidic and basic drugs can differ in where they become ionized and available for absorption.
- Acidic coatings and buffers are presented as ways to delay release of a basic drug until it reaches a less acidic part of the digestive tract.
- The example compound has a conjugate-acid pKa of 6.9; at stomach pH 1, the equilibrium calculation predicts almost complete protonation.
- At a buffered colon pH of 7.6, the ICE-table calculation gives about 17% ionization.
- The comparison shows how a fixed buffered pH changes the equilibrium result, and why the same base can have different ionization fractions in different environments.
- For a weak acid HA, the ionized form is A−, so percent ionization is calculated as [A−] divided by the total concentration of HA and A−.
- Rearranging the acid equilibrium expression produces the Henderson-Hasselbalch relationship between pH, pKa, and the ratio of conjugate base to acid.
- The lecture derives the weak-acid fraction as 1 / (1 + 10^(pKa − pH)); at pH equal to pKa, the acid is 50% ionized.
- When pH rises above an acid's pKa, the A− fraction increases, which matters because charged species tend to remain in aqueous environments.
- For a base, the protonated form BH+ is the ionized species; the lecture calculates its fraction using the conjugate acid's pKa.
- The base expression is complementary to the weak-acid fraction, so identifying whether a compound is an acid or a base is essential before choosing the equation.
- For a base with pKa 6.9, the example predicts approximately 100% ionization at pH 1 and 17% at pH 7.6.
- These results match the earlier ICE-table examples while requiring less calculation when pH is buffered.
- Naproxen is treated as an acid with pKa 4.2, while codeine is treated as a base with pKa 8.2; neutral forms favor the organic phase and charged forms favor the aqueous phase.
- At pH 2, naproxen is mostly neutral and partitions into the organic layer, while protonated codeine remains in the aqueous layer.
- Between pH 4.2 and 8.2, both compounds are mainly charged and aqueous, making that range unsuitable for separating them by phase partitioning.
- A separatory funnel lets the organic and aqueous layers be collected separately after pH adjustment and mixing.
- The example gives heroin a basic pKa of 7.9 and quinine two basic pKa values, 5.1 and 9.7.
- At pH 8.6 to 9, heroin is predominantly neutral and favors the organic phase, while quinine remains at least partly protonated and favors the aqueous phase.
- The pH window creates the needed contrast between neutral and charged forms; a separatory funnel can then collect the two compounds in different layers.
- The same pH-driven partitioning principle explains how compounds distribute between aqueous fluids and fatty tissues, and supports forensic sample cleanup before instrumental analysis.
Transcript excerpt
0:00 Okay, I'm excited about today because we're going to take what you've learned with ice tables and acid base chemistry and all of that and and apply it to forensic science and forensic chemistry. So this is a this is really I I you know I don't want you to like up and leave because you had this many students have this attitude if it's not on the test I don't want it coming in my ears right so this this is an application of what's on the test okay so just a lot of this is just should just be interesting to you and I want you to have the mindset to say oh if I'm in that field would I know how to do that like with your acidbased chemistry you would if you would think about the problem so So we're going to be talking about using Latier's principle and equilibrium chemistry to help us in forensic science. So this is using chemical
0:48 fundamentals like the pH and so on in forensic applications. So why do we want to do this? Well, we get a lot of crummy samples out in the field and bring them into the lab. And we can't just take that that nasty, you know, messed up field sample and throw it onto our really expensive and fine-tuned chemical equipment. Okay? We need to purify things. We need to separate things. We need to take um you know, in the best case, let's say we are doing a meth lab, right? And we've got solvents and the solvents are fairly pure. So that we might be able to dilute and shoot right on the instrument, but that's pretty rare. A lot of times what we see in in the field is just just nasty. Um we can
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