Fluorescent stains like EtBr, DAPI, and SYBR stains, can bind to nucleic acids (DNA & RNA), steal UV light we can’t see and give us back visible light we can see, allowing us to, for example, visualize where DNA pieces are in an agarose gel. Well, DNA & RNA can also steal UV light (without help) – but they don’t give it back! Therefore, we can use a technique called spectroscopy to measure the *absorbance* of UV light to get information about the quantity and purity of nucleic acids in solutions. We frequently do this using a device called a spectrometer, such as a “NanoDrop.” Since you’re not separating the nucleic acids by size like you do with agarose gels, you don’t get information about size, but you get information about purity and quantity (with the help of Beer’s Law)which you can’t get from a gel. Here’s what there is to tell…

video added 10/28/21

We often use spectroscopy (such as with a NanoDrop spectrophotometer) to measure concentrations of molecules like nucleic acids (DNA or RNA) or proteins – by shining light through solutions containing them and seeing how much of the light gets stolen along the way, we can calculate how stuffed full of those molecules the solution was. 

This is possible because different molecules have different tendencies to absorb different wavelengths of light to different extents (more in a sec). If you shine a molecule’s “favorite” wavelength at it, the more molecules there are, the more that will get stolen. You could just measure the absorbance of this favorite wavelength and, if you have a pure sample and you know how much your molecule likes that wavelength (the extinction coefficient) you can calculate the concentration (using a law called Beer’s law we’ll get to) just with using that one wavelength of light. 

But is your sample really pure? If other molecules like that wavelength too they can be artificially inflating your concentration – not to mention other problems with having contaminating molecules around that can interfere with things. 

So (if possible) you shouldn’t just look at a single wavelength absorption value – instead, you can learn a lot more if you measure over the whole ultraviolet (UV) and visible light portions of the electromagnetic radiation (EMR) spectrum. Each wavelength that gets absorbed will show up as a peak an absorbance spectrograph, which shows how much light of a range of wavelengths was prevented from making it all the way through the sample to the detector on the other side. 

The height of the peak you usually focus on for a particular molecule is where it absorbs best -> this corresponds to how much stuff is there (concentration) (taking into account that molecule’s willingness to absorb that wavelength (its extinction coefficient). BUT a lot of information about purity can be gained by looking at where they *shouldn’t* be absorbing much light – and ratios between peaks can tell you about how pure that stuff is

To illuminate why this is, we have to first make sure we’re clear on what “light” is – so a quick overview… 

Light (electromagnetic radiation) (EMR) can be thought of as little packets of energy called photons traveling in waves. Different colors have different wavelengths of light with different photon energies (this is also true for “invisible” colors – wavelengths outside of the visible spectrum – like radio waves, which are lower frequency and ultraviolet (UV) waves which are higher frequency. Frequency refers to how frequently the peaks of the waves come and it’s directly related to the energy – more energy, higher frequency and, since all EMR travels at the same linear speed, those peaks have to come closer together so energetic ones don’t outpace the laggards. (don’t confuse these peaks with the spectrograph peaks).

As I promised to tell you more about, different molecules absorb different wavelengths of light to different extents, and this can be quantified by a number called the extinction coefficient, which tells you how well a molecule absorbs light of a particular wavelength. Even more on *why* here: http://bit.ly/lightleafcolor 

But, basically molecules are made up of atoms (these are the individual “C”s & “H”s & “O”s in chemical structures) and atoms are made up of a dense central nucleus where positively-charged protons and neutral neutrons hang out. And negatively-charged electrons whizz around them in an “electron cloud.” Those electrons are reigned in by the opposite-charge-ness of the protons in the nucleus. I like to think of the nucleus as being like a dog walker walking a bunch of energetic electron dogs. But it’s harder to feel that charge when you’re further away so, “not feeling the love” so much, the outermost electrons are the least loyal & the most energetic (they don’t have to “waste” as much energy resisting that pull). So they’re most likely to interact with other molecules. We call these energetic, outermost electrons “valence electrons” and they’re the ones we care most about most of the time. They’re the ones that can do things like merge homes with electrons from other atoms to give you strong covalent bonds or – as is important for this case – they can absorb photons. 

They *can* but that doesn’t mean they *will* – whether or not a molecule absorbs a certain wavelength of light has to do with how much energy its valence electrons currently have and how much more energy they need to “get to the next level” – I talked about there being an electron cloud, but that cloud is like a housing district – there are places electrons like to hang out more, so, even though they’re constantly zipping around, you have the greatest chance of finding them in “electron houses” more formally referred to as “molecular orbitals”

You can think of it kinda like electrons needing to “pay rent” in the form of energy to live in various houses. And, thanks to that proton pulling power stuff I was talking about, orbitals farther from the nuclei of the molecules – out in the “electronic suburbs” – require electrons to pay higher rent (have higher energy) than those closer in. Electrons will usually live in a “ground state” where they’re as far away as they can *comfortably* get. But the housing situation isn’t set in stone – it’s possible to “upgrade” – but you have to be able to pay the difference in rent. In “exact” change. And this energy money can come from photons of light – but that light has to have the molecule’s optimal energy. 

Thanks to their unique chemical makeups, different molecules have electron houses that are spaced differently with different energy differences between them, so the amount of energy required to get excited varies from molecule to molecule, and, to get that energy from light, you need light with photons that give you “exact change” to pay the rent difference. So different molecules absorb different wavelengths of light. 

If you prefer the dog analogy, on the subatomic level it’s kinda like a dog seeing a squirrel, getting excited, and stretching its leash to get further away. different molecules have electrons that get excited by “different squirrels”

Squirrel or suburbs, an excited electron can’t stay in this outer orbital for long and will fall back down to its more cozy ground state, giving back the energy it had absorbed.  In fluorescence the absorbed energy is released as another photon (but of a different, lower energy, wavelength) – but most of the time when a molecule absorbs light it just “steals it” from the electromagnetic radiation (EMR) spectrum and, instead of giving us light, it loses the energy as heat, etc. 

But we can still detect that light was stolen –  if it steals light from the visible light part of the spectrum, we can tell something happened with our naked eyes. White light is made up of all colors of the rainbow and, if one color’s stolen (absorbed) we see the “leftovers” as a different color. A color wheel helps you see what color something will appear if the color opposite it gets absorbed.

But molecules can also absorb light we can’t see – and often this involves UV! Even though we can’t see UV, a spectrophotometer can. So we can look to a spectrograph and hope we get the last laugh. 

Don’t freak out if you see multiple and/or broad peaks – Because molecules can have lots of atoms with lots of electrons and because there’s some “wiggle room” with regards to energy levels, molecules don’t just absorb a single wavelength. They often have a wavelength they absorb the most strongly at, corresponding to the electron most readily promoted, with some lower absorbance on either side (like a bell curve) corresponding to the wiggle-roominess (things like vibrational levels in between the main energy states) and then possibly some secondary peaks elsewhere corresponding to different electrons being promoted. 

You can see this if you look at a UV-Vis absorbance spectrum, which you can can get if you use a spectrophotometer like a NanoDrop. To get a full absorbance spectrum, you shine light of all wavelengths (well, at least all visible & some ultraviolet (UV), which is where biological molecules tend to absorb) & measure what goes through (is transmitted). Whatever doesn’t go through is assumed to be absorbed (or abducted by aliens…). So absorbance = 1-transmission.

So how does it work? We take whatever solution we want to measure and put it in a plastic or glass holder called a cuvette which has a window for light to shine through & stick it in a spectrophotometer which actually shines the light through one side & measures what comes out the other. Different cuvettes can have different path distances the light has to travel, and the Beer Lambert law takes this into account – the longer the path, the more molecules light’s likely to hit (regardless of the concentration) – and the more molecules it hits, the more chances there are to be absorbed. Our calculation of concentration is based on how much light gets absorbed so we need to account for this distance.

It’s also important to have a blank – this is just the liquid you have your samples in and all the “constant” components (e.g. salts, etc. that are in each reaction) – its made up of molecules too so it’ll have a characteristic absorbance that will always be there, whether or not the reaction we’re looking for actually occurs. And some of its absorbance spectrum might overlap with our products’. So we want to subtract it out so we don’t confuse it for our signal.

Cuvettes are great for things where you have “large” samples. But if you don’t have much sample though, you’ll want something smaller-scale.

The NanoDrop spectrophotometer has a little pedestal you put a drop of liquid on  (a really tiny drop, like 1-2μL “μL” stands for microliter and it’s a millionth-of a liter, or a thousandth of a milliliter. Then you lower the arm -> it contacts the liquid then pulls up a little bit and, when it does, it pulls on the liquid. It does this thanks to surface tension. Surface tension occurs because the molecules of the liquid like each other more than they like the air – so they try to stay together & maximize the liquid-liquid interactions while minimizing their combined air exposure. more here: http://bit.ly/surfacetensionbubbles 

When you put the drop on, surface tension causes it to remain drop-like. But when you lower the arm & squish it down, some of the water molecules stick to the top surface. And when the arm pulls back up, these molecules get lifted – and the other water molecules don’t want to leave their friends behind -> as a result a column of liquid forms

This column is just like the column of liquid in the cuvette except it’s much smaller and it’s held up by surface tension instead of being barricaded by glass/plastic. And it’s “sideways” – the light travels through from above and is recorded below

Another difference about this column is that the path length’s adjustable (the column can be pulled up higher or squished) – so it can adjust for different concentrations (e.g. if your sample’s too concentrated it’ll see this squish down to shorten the path length so the light doesn’t meet as many, or if it’s too dilute it can pull up to make sure lots of molecules get hit by light). The NanoDrop software then has to correct for this when it uses Beer’s Law to get to concentrations

A couple things we commonly use it to measure are nucleic acids and proteins. It’s easier to get columns to form for nucleic acids than proteins because the proteins often weaken the surface tension – not all parts of proteins like water, so it’s harder to get continuous water networks to stay as you pull up the column —> column breaks. But there’s strength in numbers, so, when loading protein, I usually load 2uL, but I load 1 or 1.2 for nucleic acids.

Later we’ll look at how we can use the absorption at a single wavelength to calculate the concentration, but first let’s look at what you can learn from also looking at the whole spectrum. Because molecules have overlapping spectra (e.g. both DNA & RNA absorb light of 260nm wavelength) you look to ratios. A couple key ratios for when you’re analyzing nucleic acid (DNA or RNA) quality are 260/230 & 260/280.

In terms of biomolecules, at 260 – dominant absorbers are DNA & RNA & at 280 – dominant absorber is protein

260/280: tells you about protein “contamination” – I put contamination in quotes because if you’re analyzing protein, it’s the DNA that’d be the contaminant!

260/230: tells you about “contamination” from proteins and/or things like phenol or salts that are left over from the purification process

But why? Where does that absorption come from?

DNA & RNA only have 4 letters each, and all of them absorb 260nm light, but it’s their unique part that does the absorbing – adenine, guanine, cytosine and thymine/uracil bases. These are aromatic rings (aromatic rings are rings where there’s some communal electron sharing that stabilize them – more here: http://bit.ly/phenylalaninearomatic). This electron delocalization through resonance reduces the cost to move (to promote an electron), so aromatic rings are often present in things like dyes.

Resonance is also involved in why proteins absorb light. Proteins peak at 280 & 230. The parts of proteins that absorb at 280 are aromatic rings (sound familiar?) Only 3 of the 20 common amino acids have these – Tryptophan (Trp)  & Tyrosine (Tyr) are the major contributors – Trp the most so – Phenylalanine has a ring too, but it doesn’t absorb here as much. Cysteine crosslinks can also absorb, where applicable – and different proteins have different numbers of all these. So different proteins absorb 280nm light differently, which is reflected by different extinction coefficients. Proteins also absorb at 230nm and that absorbance is from the generic backbone part – corresponds to absorbance by the peptide bonds linking the letters. These peptide bonds also have resonance, but not as much as rings do, and they absorb ~190-230nm. More on using UV to measure protein concentration here: http://bit.ly/proteinmeasuring 

DNA & RNA bases all absorb at 260, but to different extents-> If you measure the 260/280 ratios for each nucleotide separately you get: Guanine: 1.15; Adenine: 4.50; Cytosine: 1.51; Uracil: 4.00; Thymine: 1.47

one thing you might notice is that uracil (U) which is in RNA but not DNA has a much higher 260/280 than its DNA counterpart, thymine (T) – as a result, pure RNA has a higher 260/280 ratio than pure DNA

If you take the weighted average of the different bases into account, pure RNA should have an A260/A280 ratio of ~2 – it absorbs ~ 2 times more 260nm light than 280nm light) & pure double-stranded DNA, which only absorbs ~1.8 times more at 260 vs 280, should have an A260/A280 ratio of ~1.8

Because DNA absorbs so strongly at UV260, where protein doesn’t, it’s relatively easy to see if you have DNA in your protein prep, but it’s harder to tell if you have protein in your DNA prep – 260 will dominate the 260/280 ratio

You also want to look at the 260/230 values -> for pure nucleic acids, these are usually ~2.0-2.2

So far, the “contaminants” we’ve discussed are biomolecules that co-purified with the molecule you purified, but contaminants can also come from chemicals you used in the purification process. For example, phenol & chaotropic salts like guanidine are often used when purifying nucleic acids, such as with phenol-chloroform or Trizol extractions (more here: http://bit.ly/2Xj4Zyc )

These absorb strongly at 230nm, which (in addition to the fact that proteins also absorb there) is why a low 260/230 ratio could be concerning if you’re looking at a nucleic acid prep.

Now, as promised, let’s get discuss how we can convert absorbance to concentration using Beer’s Law. We can characterize how much a molecule absorbs light at any wavelength (we usually choose its “favorite” – peak absorption) using its extinction coefficient.

The equation is: A = εcl

A = absorbance

ε = extinction coefficient (aka molar absorptivity coefficient) – specific for particular molecule & particular wavelength; units of L mol⁻¹cm⁻¹

c = concentration (in mol/L) – this is molarity – a mole is just a chemist’s “baker’s dozen” – it’s Avogadro’s number (6.022 x 10²³) of something – solute molecules or donuts, it’s just a number http://bit.ly/c1v1equalsc2v2 

l = path length (in cm)

rearrange that a bit and you get 

c= A/εl

Even though the base composition will affect the exact extinction coefficient (the measure of how well a specific something absorbs a specific wavelength – the thing you stick into Beer’s law to convert absorbance to concentration, you can still estimate extinction coefficients for “generic” DNA & RNA

dsDNA: (0.02 μg/mL)^-cm

ssDNA: (0.027 μg/mL)^-cm

ssRNA: (0.025 μg/mL)^-cm

To make these easier to use, there’s a shortcut. “Standard Coefficients” which assume a 1cm pathlength (if you’re using a different one, then divide the A by that. https://bit.ly/3dg5yDi 

Basically, since 

c= A/εl, then c = A/l * 1/ε

if l = 1cm, then this is c = A * 1/ε

So if we calculate 1/ε, then we can just multiply that by A to get the concentration. So, instead of remembering 0.02, 0.027, & 0.025, remember

dsDNA: 1/0.02= 50μg/mL(ng/μL)  -> A260 of 1.0 corresponds to 50μg/mL(ng/μL) of pure dsDNA. A260 of 2.0 corresponds to 100μg/mL(ng/μL) of pure dsDNA. etc… 

ssDNA: 33μg/mL -> A260 of 1.0 corresponds to 33μg/mL(ng/μL)of pure dsDNA. A260 of 2.0 corresponds to 66uμg/mL(ng/μL) of pure ssDNA. etc… 

note: 1/0.027 is actually ~37μg/uL so I honestly don’t know what’s going on here, but most sources say 33?! 

RNA: 40μg/mL (this one makes since again since 1/0.025=40) -> A260 of 1.0 corresponds to 40μg/mL(ng/μL) of pure dsDNA. A260 of 2.0 corresponds to 80μg/mL(ng/μL) of pure ssRNA. etc… 

From these values you can see that strandedness matters. When DNA is double stranded, its bases (the parts that absorb the light) are “hidden” & “busy” binding to the base on the other strand. But with single stranded DNA, the bases are out in the open and free to absorb. As a result you get a hyperchromic shift -> single-stranded DNA (ssDNA) absorbs more UV light than double-stranded DNA (dsDNA) & free nucleotides absorb more strongly than either of those.

final note: You also want to look at the absorbance at 320nm. This tells you about the turbidity of your sample – if your sample is “cloudy” – suspended particles in the solution can scatter the light, preventing it from reaching all your molecules. So this gets adjusted for. So for DNA, you can get concentration using this equation: concentration (ug/mL) = (A260 reading – A320 reading) x dilution factor x 50ug/mL

more on fluorescent DNA stains: http://bit.ly/fluorescentstains 

suggested reading: What Causes Molecules to Absorb UV and Visible Light. (2020, August 21). Truro School in Cornwall. https://chem.libretexts.org/@go/page/3746

more on topics mentioned (& others) #366DaysOfScience All (with topics listed) 👉 http://bit.ly/2OllAB0

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