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Introduction to buffers | Water, acids, and bases | Biology | Khan Academy

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- [Voiceover] In order for our bodies to function properly, the pH of our blood has to be within a fairly narrow range. Our pH, so the pH of blood needs to be between seven point three five and seven point four five. If the pH falls below seven point three five, you're going to be diagnosed, and you know this line is, there is kind of a gray area here but this is what's defined by the medical community if your pH falls below seven point three five, you're actually diagnosed with acidosis. Your blood is becoming too acidic. You have too high of a hydrogen ion concentration in your blood, or a hydronium concentration in your blood. If your pH, if the pH in your blood gets above seven point four five, you're diagnosed with alkalosis. Alkalosis. Your blood is too alkaline. It's becoming too basic. The hydrogen concentration, hydrogen ion concentration in your blood is getting too low. And so you might say, "Wow, you know this feels like a fairly narrow range. How does blood, how does our body put up with acidic things, acidic molecules entering our blood, or basic molecules entering our blood. How can it handle that without our pH, while keeping our pH in this range right over here?" And the answer lies in something that's also useful for the transportation of carbon dioxide in our blood. And this right over here, these equilibrium reactions, we see that carbon dioxide, when you put it in the blood, which is primarily water, so carbon dioxide in aqueous solution, it will react. And we have some enzymes that help this along, but it will react to form carbonic acid. Let me write this down. It will react to form carbonic acid, which is a weak acid, this is carbonic adic. And then that can dissociate to form bicarbonate. Let me write that over here. Bicarbonate. Bicarbonate. And a hydrogen ion, or when we know that that would just attach to a water molecule and become a hydronium ion. And so, why is this useful? Well, I said, it's actually part of the carbon dioxide transportation in the blood because, based on the sources I've looked at, about five to 10 percent of your carbon dioxide can just dissolve in the blood and then also another, roughly, five or 10 percent can actually be bound to hemoglobin and be transported that way. But the bulk of it actually needs to go through, needs to actually be transformed, needs to react to become carbonic acid and bicarbonate in order to be transported. In fact, in your blood, most of the carbon dioxide in your blood is actually in this form right over here. And in particular, bicarbonate. In the sources that I look at, 80 to 90 percent of the carbon dioxide in your blood is being transported in these forms, and primarily as bicarbonate. So this isn't the topic of this video, what's a useful way to transfer, to transport carbon dioxide in your blood, but this is how we do it. But the topic of this video is why this is also useful for maintaining our blood pH in this range. Because these equilibrium reactions between carbon dioxide, carbonic acid, and bicarbonate this is a buffer system. This is a buffer, this is a buffer system. And the word "buffer," in our everyday language, it refers to something that kind of smooths the impact of something, or it reduces the shock of something. And that's exactly what's happening here. Let's think about, remember, these are all equilibrium reactions, this is a weak acid, and you can even look at the different constituents of these molecules and account for them. You have one carbon here, one carbon here, one carbon there. You have one, two, three oxygens there. You have one, two, three oxygens there. One, two, three oxygens there. You have two hydrogens, two hydrogens, two hydrogens. But let's just think about what if you started dumping hydrogen ions in the blood. So, what if you were to dump hydrogen ions, what's going to happen? Well, if you dump more hydrogen ions, if this right over here increases. Actually, let me put it this way, if you were to just dump hydrogen ions and if you didn't have this buffer system, then your pH would decrease. Your pH would go down, and if you do it enough, your pH, you would end up with acidosis. But lucky for us, we have this buffer system. And so if you increase your hydrogen ion concentrations, Le Chatelier's principle tells us, "Hey, these equilibrium reactions are going to move to the left." So the more hydrogen ions you have sitting in the blood, the more likely they're gonna bump into the bicarbonate in just the right way to form carbonic acid. And the more carbonic acid that you have in the blood, well, it's the less likely that you're going to have the carbon dioxide reacting with the water to form more carbonic acid. So, as you add more hydrogen ions, they're just going to be sopped up by the bicarbonate. So this equilibrium, this set of equilibrium reactions is going to move to the left. So you're not going to have as big effect on pH. And similarly, if you dumped some base, let's say, you dumped some base in your blood right over here, well, instead of it just making your pH go up, and possibly give you alkalosis, well now, the base is going to sop up the hydrogen ions, and typically that would just make your pH go up, but if you have these things going down, well then, you have fewer of these to react and have the equilibrium reaction go to the left and so the reaction is going to move more and more to the right. And so this reaction, you're just gonna have more carbon dioxide being converted to carbonic acid being converted to bicarbonate. This whole thing is going to move to the right. And so it's going to be able, to some degree, replace the lost hydrogen ions. So this right over here's a buffer system. It helps dampen the impact, as if you have more hydrogen ions enter the system, or as if you have something sopping up all of the hydrogen ions. And it's super important for us, well, just being able to live. And frankly, all mammalian systems.

Video Details

Duration: 6 minutes and 20 seconds
Language: English
Genre: None
Views: 12
Posted by: jake.goldwasser on Nov 19, 2015

Introduction to pH and the pH scale. Examples of calculating pH of pure water, bleach, and orange juice.
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