Buffer Solution

in this video, we're going to review buffer solutions. So recall from your previous chemistry courses that buffers are substances that resist changes to P. H. And this is true even when small to moderate amounts of a strong acid or base are added. And so, in other words, when small to moderate amounts of a strong acid or base are added to a buffer solution, the pH of that buffer solution is on Lee going to change a little bit in comparison to the pH change that we would see if the same amount of a strong acid or base were added to an UNB, uh, furred solution. And so down below. In our example, we have four different scenarios numbered 123 and four. And in the first two scenarios over here on the left, we're adding a small amount just one mil leader off a strong acid hydrochloric acid with 0.1 Moeller concentration. And so, in the first scenario over here noticed that the solution in our early Meyer flask is an UNB, uh, furred solution. And so when we add this little bit of a strong acid to the UNB, uh, furred solution notice that the pH drops it goes much, much lower, and so the pH indicator that's inside of the solution changes color. It changes from a blue color to a pink color down below to show the change in the pH that we see, uh, in this unbuttered solution. Now on the right. We're here in Scenario two, where we have a buffered solution when we add the same exact amount, the same exact small amount of a strong acid Notice that the pH remains relatively stable and the pH indicator does not change color. So it has a blue solution up above, and it remains blue to show that the pH remains relatively stable and does not change so again. The point here is to show that it's able. The buffered solution is able to resist changes in pH now in scenarios three and four. In a similar way, we're adding a small amount of this time a strong base sodium hydroxide with a concentration of 0.1 Mueller. And so in scenario three, we haven't unbuttered solution, and when we add that small amount of a strong base to the unbuttered solution, notice that the pH goes much, much higher. And that's because the solution is UNB, uh, furred. And so what you'll notice is that there's a color change from, ah, blue color up above to a purplish color down below to show that the pH changed. Now again with the buffered solution, when we add that same small amount of a strong base to the buffered solution, noticed that the pH remains relatively stable. So we have the same color change or no color change. Light blew up above and light blue down below. And so again, the major take away here is that buffers can resist. Changes in pH. Now over here on the right were reminding you guys that the Henderson hostile back equation can actually be used to prepare buffer solutions. And so recall from our previous lesson videos that Henderson Hasselbach equation is expressed as the pH of the solution equal to the PKK of an acid, plus the log of the final concentration of constant conjugate base over the final concentration of conjugal Aston. And so in our next practice, videos will be able to practice utilizing the Henderson Hustle back equation while preparing buffer solutions. So I'll see you guys in that practice video

A. What volume of 0.1 M acetic acid (pK _a = 4.8) is required to make 1 liter of 0.1 M buffer solution at pH = 5.8?

B. What volume of 0.1 M sodium acetate is required to make the same buffer solution?

a. 193 mL                                                       

b. 91 mL 

c. 909 mL 

d. 807 mL

• A. 193 mL
• B. 91 mL
• C. 909 mL
• D. 807 mL

so the most effective buffers are actually weak acids and bases, and their congregates and living systems tend to use weak acids as their buffers. And the effect of buffering range of a weak acid is centered right around the inflection point or the midpoint of a tight trey shin curve, which means that the effect of buffering range is centered right around the PK. And really, the effect of buffering ranges within one unit of the peak. A. So what this means is that if the pH of the solution is within one unit of the PKK than this substance is capable of acting as an effective buffer. So let's take a look at an example, and we're gonna look at the effect of buffering range of acetic acid. And as you guys already know, acetic acid has a chemical formula of ch three C 00 H. And acetic acid is a mon, a product weak acid, which means that it only has one acidic hydrogen, which is shown in red here, and the PKK of this acidic hydrogen is and so when I see the acid loses its hydrogen, it becomes an acetate conjugate base And so the acetic acid buffering range, effective buffering range is going to be within plus or minus one of the P K. So if we subtract one from the peak A. What we get is 3.8. And if we add one to the PK, what we get is 5.8, So the effect of buffering range of acetic acid is from 3.8 to 5.8. And let's take a look at the Thai Trish in curve of acetic acid with a strong base and again toe orient you guys on the Y axis, we have the pH of the Ana light solution, and on the X axis we have the amount of Thai trahant that's being added. And so notice here that what we have is our inflection point right at the 0.5 Mueller equivalents added, which means that the inflection point or the midpoint is right here, and that corresponds with the pH of 4.8, which means that the P K is equal to 4.8, and the effect of buffering range, we know is within plus or minus one of the P K. So that means it's gonna be from 3.8 up to 5.8 or in this pink range here. And that means that this entire section of the, uh, titrate shin curve here represents the effective buffering range of acetic acid. And so what this means is that when the pH of the solution falls within this range here, then acetic acid is capable of acting as an effective buffer. And so we're going to get some more practice with the effective buffering ranges in our practice problems, so I'll see you guys in those videos.

Which of the following compounds would make for the best buffer at pH 8?

• A. Acetic acid, pKa = 4.8
• B. Tricine, pKa = 8.15
• C. Glycine, pKa = 9.9
• D. Tris, pKa = 8.3

so buffers air actually supercritical toe life and most biochemical processes require enzymes and enzymes require a very specific pH in order for them to work properly. And so living systems need to make sure that they have a way to maintain that. PH to make sure that they're enzymes are working properly and to make sure that their biochemical processes actually occur. And so living systems tend to use weak acids as their buffers and as a way to maintain home yo Stasis by maintaining their pH. And so some buffers maintain intracellular pH, whereas other buffers maintain extra cellular pH or pH on the outside of cells. And so the phosphate buffer system, which uses di hydrogen phosphate as well as hydrogen phosphate, maintains intracellular pH, whereas the bicarbonate buffer system maintains extra cellular pH. And so let's take a look at our example below and notice in green. Here we have our cell, and we have our plasma membrane, which separates the inside of the cell from the outside of the cell and inside of the cell. We have our phosphate buffer system, which maintains intracellular pH and on the outside of the cell we have our bicarbonate buffer system, maintaining extra cellular pH and so the phosphate buffer system. We know that phosphate groups were found on so many different types of macro molecules, including nucleotides, which each have a phosphate group and also lipids such as fossil lipids. And so here we have a di hydrogen phosphate and the PK is right around seven to, which means that the effect of buffering range is gonna be within plus or minus one of 7.2. So it'll be from 6.2 toe 8.2, and most living systems have an intracellular ph of right around seven. And so seven falls right within this effective buffering range, which means that di hydrogen phosphate and hydrogen phosphate are capable of acting as effective buffers within the cell now on the outside and sell. We have the bicarbonate buffer system and recall from your previous courses that the bicarbonate buffer system is used to maintain the pH of blood and the extra cellular fluid, and so the bicarbonate buffer system consists of carbonic acid and bicarbonate and carbonic acid has a P K A. Of 6.4, which means that the effect of buffering range is from 5.4 up to 7.4. And so, uh, the pH of blood is maintained right around 7.4, and that falls right on the edge of the effective buffering range. So you might be thinking this is not an effective system to maintain, uh, the pH of blood. However, there's an additional component here that will talk more about later in our course. And that's this exchange of CO two with the environment. So by breathing out CO two and we can control the amount of CO two this present and so therefore we can control and extend this buffering range even further and again, we'll talk more about that later in our course. So for now, this is a good summary of the biological buffers, and I'll see you guys in our practice videos.

MOPS (pK_a =7.2) is a weak acid & acts as a buffer. Calculate the ratio of its basic/acidic species at pH = 6.0.

• A. 0.098
• B. 1.24
• C. 0.377
• D. 0.063