You’re preparing buffers in the lab, and maybe you’ve been following a protocol for the longest time – just going through the motions. Now, though, it’s time for you to prepare and optimize your own buffer. What’s actually involved? What are the steps and what is the process of preparing buffers for your experiments?
Buffer preparation spans six steps – selecting your buffer, deciding what chemicals to use and in what amounts, measuring them out, mixing them, adjusting the pH, and making up the right volume. All of these steps require careful optimization to ensure experimental success.
Here, we will look at the basic steps involved when preparing any buffer, the common mistakes you may encounter and the best practices to follow to avoid those mistakes.
Specifically, we will take a glance at chemical calculations and handy formulas, best practices to follow, and what’s involved when adjusting the pH and the final volume.
After that, we will also discuss some important things you need to keep in mind while storing and using the buffers you make.
Article table of contents:
Once you know which buffer you need for your experiment, the preparation steps involve checking the recipe, calculating the amounts of the different chemical constituents, measuring them out, mixing or dissolving them, adjusting the pH, and making up the desired volume.
Here is a brief description of each of these steps, including the things you need to follow as best practices and the common mistakes you want to avoid.
The buffer you need is largely based on the pH required for your experiment. The pK value of a buffer gives the range of its buffering capacity, or the pH range it can maintain.
Also, there are buffers that are standardized for certain experiments – for example, TAE for DNA gel electrophoresis.
So, you would want to choose a buffer standardized and established for what you want to do in the experiment. If you are not sure about which buffer you want – here is a general rule: Check the pK of common buffers, and then choose one for which the pK is within +/- 1 of the pH you want to maintain during your experiment.
Once you determine which buffer you need, check for the recipe.
Preparing buffers largely depends on the basic concepts of chemical calculations– Avogadro’s number, molecular weight, mole, chemical equivalence, molarity, normality and molality.
Let’s imagine you are used to working with a standard buffer solution recipe in your lab, but now your experiment calls for a different concentration. Maybe your current recipe makes a 20 mg/ml solution, but you need to make a 10 mg/ml solution.
These types of situations are so common and they are why chemical calculations are very important.
So, here is a very quick recap of those terms and calculations while not getting into the intricacies of these concepts.
Mole is the amount of a substance – an element or compound that has Avogadro’s number of molecules of it. And Avogadro’s number is 6.022*1023.
NaCl powder, for example, corresponds to the molecular weight in grams – also called gram molecular weight.
So, in this case, for NaCl, 1 mole means 58.44 g – this weight of pure NaCl has 6.022*1023 molecules of it.
For a gas, where weighing it is very difficult – practically impossible in a regular lab without sophisticated instruments, 1 mole corresponds to 22.4 liters (L) of that gas. So, 1 mole of CO2 is 22.4 L of that gas and 44 g (molecular weight of CO2) of it in weight.
Molarity is very commonly used to indicate concentration of solutions. Technically, molarity or molar concentration, represented as M, is a measure of the number of moles of a substance in a liter of its solution. 1M means 1 mole of the substance per liter of the solution; 2M means 2 moles of the substance per liter of the solution and so on.
So, going back to the example of NaCl, 1M solution of NaCl in water means 58.44 g of NaCl dissolved for liter of the solution.
Molality, also called molal concentration, is the number of moles of a substance per kilogram of the solvent. So, for 1 molal NaCl solution in water, 58.44 g of NaCl are dissolved in 1 kg of water.
Molality is less commonly used as a measure of solution concentrations compared to molarity. However, it is something to keep in mind just in case you have a solution where the solute concentration is indicated in molality.
For water as a solvent, the amount in 1 kg is equal to 1 L because the density of water at room temperature is ~1000grams per liter. So 1 kg and 1 L of water basically mean the same.
1molar NaCl in water and 1molal NaCl in water are same – both mean 58.44 g of NaCl dissolved in 1 L (or 1 kg) of water. However, it’s not always a one-for one conversion. When the density of a 1 kg is not equal to 1 L (oils for example), then 1molar is not equal to 1molal.
Another thing to note is for a simple solution such as NaCl in water, the volume of 58.44 g of NaCl is negligible.
However, for complicated solutions that you commonly need to make for bioscience experiments, there are so many solutes, that if you add all of them together in a liter of water, you might end up having a volume more than 1 L.
As an example, suppose you have three solutes A, B and C, with molecular weights of 100 g, 200 g and 300 g respectively. You want to mix A, B and C together to have a solution that is 1molal for A, B and C. This is straightforward – you measure 100 g of A, 200 g of B and 300 g of C and add to 1 L of water and mix. That gives you a solution that is 1molal for A, B and C.
But, when you measure the total volume, you find that it is more than 1L. The addition of such bulk amounts of solutes increased the overall volume. So, the solution you made is not 1molar for A, B and C because the total volume is more than a liter. The way around this is making up the volume – a concept we will discuss later in this article.
Normality, represented as N, is defined as the number of mole equivalents, or gram equivalents, of a substance dissolved per liter of the solution. In other words, 1N of a solute means 1 g equivalent of it dissolved per liter of the solution.
The concept of chemical equivalence, or equivalent weight, of any element or compound is complicated, and we will not get into the chemistry intricacies. For the purposes of making solutions for your experiments in the lab, here are some handy points:
For a salt, the equivalent weight is molecular weight divided by the valency of the cation. For example, equivalent weight of NaCl is 58.44/1, or 58.44 g. So, 1M solution of NaCl in water is the same as 1N NaCl.
However, for CaCl2, whose molecular weight is 111 g, the equivalent weight is 111/ 2 (because Ca is divalent here), or 55.5 g.
So, for CaCl2, 1M solution means 111 g of CaCl2 in 1 L of solution, while 1N means 55.5 g CaCl2, dissolved in a liter of the solution. In other words, 1M CaCl2 is the same as (or, is equivalent to, in chemical terms) 2N CaCl2.
For an acid, the equivalent weight is molecular weight divided by the number of replaceable hydrogen atoms.
Basically, it is the number of H+ ions the acid yields in a solution of water. For HCl, it is 1; for H2SO4, it is 2. So, 1M HCl and 1N HCl mean the same. But 1M of H2SO4 is 2N.
Very often, you will encounter situations in lab while making solutions. For instance, you have a stock solution of 2N NaCl, and you need to make 500 ml of 0.3N NaCl. Here is a handy formula to use: V1 (in ml)* S1 (strength in Normality) = V2 (in ml)* S2 (strength in Normality). It’s very similar to the classic (C1)(V1) = (C2)(V2) formula.
So, in this example, you need to make 500 ml of NaCl with a strength of 0.3N. The volume you need to add from the 2N stock solution can be derived as:
(V1)(2N NaCl) = (500 ml)(0.3N NaCl)
2N NaCl) =150 (ml* N)
2N NaCl) = (2 N NaCl)
(V1) = 75 ml
So, you take 75 ml of the 2N stock solution of NaCl in a beaker, and make up the volume to 500ml by adding 425ml of water to it.
This formula, V1*S1=V2*S2 can be used when the strengths are expressed in molarity as well. But remember, the units of volume and strength should be the same on both sides of the equation, meaning that if you’re working with molarity instead, your S1 and S2 (or C1 and C2) must both be expressed in molarity. You cannot mix normality and molarity in the same equation.
Measuring out chemicals when making solutions involves measuring the weights of solid reagents and volumes of liquid reagents. This step needs to be done very carefully; otherwise this may affect the composition of the solution you are making.
Here are some best practices for measuring out liquid chemicals:
- For small volumes, use a pipette instead of a beaker.
- For large volumes (generally more than 5ml), measure out chemicals using a beaker kept on a flat surface such as the counter top.
- When measuring using a beaker, measure the middle of the meniscus at eye-level. Please see the figure below as a reference.
- Make sure that the beaker does not react with the chemical you are measuring. While plastic beakers are fine for most chemicals, strong acids may corrode plastic beakers. Glass beakers are a better option in these cases.
- For liquid chemicals with a high viscosity, such as glycerol, take care to ensure that the entire volume of chemical leaves the walls of the beaker or pipette tip when you are dispensing it.
Pictured is a graduated cylinder filled
with liquid. The arrow points to the meniscus where liquid measurement should
For weighing solid chemicals, here are some tips:
- Use an appropriate weighing machine. For example, if the required amount is 1 mg, obviously you would not want to measure using a balance where the margin for error itself is 0.1 mg.
- Measure the chemical within the glass chamber of the machine to avoid environmental contamination such as moisture. This is especially pertinent when dealing with a hygroscopic chemical.
- Always add the chemical slowly to the weigh boat because in the case of excess, any amount that must be subtracted cannot be put back into the container.
After measuring the chemicals out, they are mixed in a small volume of the solvent (water in most cases), and then the pH is checked and adjusted to the desired value, and then the volume is made up to the desired pre-determined volume.
To elaborate on this, here is an example.
Let’s say you have to make a 500 ml solution that is 1M for salt A and 2M for salt B. The pH needs to be 7.
For this, you make the required calculations and make the 500 ml solution. However, when you check the pH, you find it is 7.9.
So, you would need to add an acid, such as HCl, to bring the pH down to 7. But, as you add the HCl to the 500 ml solution, the volume obviously increases to more than 500 ml.
Now you have a solution that has the correct pH of 7, but the molar concentrations of salts A and B have gone below the required values of 1M and 2M since the volume is now more than 500 ml.
The way around this is to first adjust the pH and then make up the volume. So, here is what you should have done in our example.
You calculate the required amounts of salts A and B for 500 ml to achieve the required molar concentrations, but you dissolve them in only 300 ml of water. Now you check the pH and adjust it by adding HCl.
Suppose you had to add 20 ml of 1N HCl. Now, the total volume is 320 ml – 300 ml to start with plus 20 ml of HCl. Then you make up the volume to 500 ml by adding 180 ml water.
If it is pure distilled water that you are adding, it should not alter the pH at all. Pure distilled water does not have salts/ acids/ bases dissolved in it; so adding pure distilled water to make up the volume will not change the pH.
Figure 2. The wrong (pink) and right (green) approach to adjust the pH and making up the volume of the solution you are preparing.
Here are some specific tips for you to remember:
- Always test the pH and adjust it first and then make up the required volume. The initial volume of solvent should be no more than 60-70% of the total desired volume.
- When adjusting the pH, use a strong stock of acid or base depending on whether you want to reduce or increase the pH of your solution respectively. Using a strong acid/ base ensures that a low volume, instead of a large volume of it would be required to adjust the pH.
- Suppose you are adding NaOH to raise the pH of your solution to 8 from 10. Since the original composition of your solution does not have this base (NaOH), and you are adding it just to adjust the pH, you would want to keep the volume added as low as possible. Using a strong NaOH solution means, you would have to use a very low volume of it to get the pH up to 10. If the NaOH you’re using was weak, you’d need to use a much higher volume. And you don’t want that because NaOH was not in the original recipe of your solution.
- Keep stirring your solution continuously while adjusting the pH. This ensures that the acid or base you are adding gets thoroughly mixed in the solution and you are not getting the reading of the local area of your solution where the bulb of the pH meter is placed.
Figure 3. Illustrations of pH readings in different areas of a beaker. The frame on the left (pink) shows how pH will vary in different locations when material isn’t evenly distributed. The right frame (green) shows even distribution with a more correct pH reading.
After you prepare the buffer, it needs to be stored properly so that it is not contaminated and that the constituents are not degraded or lose their chemical or biological properties.
Here are quick tips:
- Store the buffer at a temperature that is suitable for its most thermolabile component. In other words, if you have a buffer with constituents A, B and C – where A and B are stable at room temperature, but C needs to be stored at 40C, this buffer needs to be stored at 40C.
- When using a stock buffer, check that no precipitates have formed. If you do see precipitate, resuspend them by using stir bars.
- Store the buffer in an inert material. Most commonly borosilicate glass bottles are preferred in case the buffer has any component that may be corrosive/ reactive to plastic. However, glass bottles are prone to breakage and are also relatively more expensive than plastic bottles. So, if the buffer is safe to be stored in plastic, glass bottles are avoided to avoid breakage.
- Be as careful as possible that no contaminants get into the buffer. For this, a best practice is to dispense exactly whatever volume you need – nothing that comes out of the container must go back in, lest that creates contamination.
While the steps of buffer preparation and storage are fairly straight forward, there is a lot of precision involved. How you exactly prepare your buffer, and any specific things you’d need to keep in mind, would depend on the exact buffer you are making. For that, here is an article where we discuss buffer categories and some specific that would get you more granular information – building on from the generic stuff that we covered here regarding buffer preparation and storage.