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Introduction to buffers

Introduction to buffer systems, which regulate pH in blood.

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  • piceratops ultimate style avatar for user Apoorva  Doshi
    What if CO2 conc. rises?
    then the buffer wont work, right?
    (28 votes)
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    • spunky sam blue style avatar for user Ernest Zinck
      Yes, the pH of the blood is controlled by the bicarbonate buffer system:
      CO₂(g) + H₂O(l) ⇌ H₂CO₃(aq) ⇌ H⁺(aq) + HCO₃⁻(aq)
      If the concentration of CO₂ temporarily gets too high, the ability of the buffer to control pH may be temporarily overloaded.
      Fortunately, too much CO₂ in the blood triggers a reflex that increases breathing.
      The blood flows through the thin membranes in our lungs, so the excess CO₂ can easily diffuse out of the blood through these membranes and be expelled into the atmosphere when we exhale.
      This loss of CO₂ shifts the position of equilibrium to the left, and the normal pH of the blood is restored.
      (69 votes)
  • blobby green style avatar for user vanya seelam
    What is the written definition of the buffer system?
    Thank You!
    (14 votes)
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  • leaf red style avatar for user nathan.birks
    What are the symptoms of acidosis and alkalosis?
    (9 votes)
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    • starky sapling style avatar for user Karen Richards
      I was curious, too! I went here for details: https://labtestsonline.org/understanding/conditions/acidosis
      At that link to the American Association for Clinical Chemistry, this is said:

      The lungs and kidneys are the major organs involved in regulating blood pH.

      •The lungs flush acid out of the body by exhaling CO2. Raising and lowering the respiratory rate alters the amount of CO2 that is breathed out, and this can affect blood pH within minutes.

      •The kidneys excrete acids in the urine, and they regulate the concentration of bicarbonate (HCO3-, a base) in blood. Acid-base changes due to increases or decreases in HCO3- concentration occur more slowly than changes in CO2, taking hours or days.

      :
      (28 votes)
  • male robot hal style avatar for user Dovid Shaw
    At , Sal explains that when you dump (H^+) into the blood, the buffer system prevents the pH from going up by creating an equilibrium that bonds the H+ to the Bicarbonate. Why does that prevent the pH from going up? There is still a larger amount of H+ in the blood.
    (8 votes)
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  • purple pi purple style avatar for user Jonathan Locke
    Is it possible to give yourself acidosis by drinking too much orange juice (or similar acidic substance) or is the buffer system too strong to be affected by what we choose to eat and drink?
    (9 votes)
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  • marcimus orange style avatar for user S Chung
    Sal mentions that when H+ increases in blood plasma, the buffer system will regain acid-base homeostasis by reacting bicarbonate with the additional H+, which in turn, forms carbonic acid. Given the increase in carbonic acid, carbon dioxide is less likely to react with water to form carbonic acid.

    While this all makes sense, what happens to the carbon dioxide that remains in blood plasma? I can't imagine that the influx of carbon dioxide would slow down on account of an increase in carbonic acid. I mean, carbon dioxide is always present, so how else is it transported if only 5 - 10% is transported by means other than carbonic acid/bicarbonate?
    (3 votes)
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    • spunky sam blue style avatar for user Ernest Zinck
      We exhale the excess CO₂.
      We have the four interconnected equilibria.
      CO₂(g) ⇌ CO₂(aq)
      CO₂(aq) + H₂O(l) ⇌ H₂CO₃(aq)
      H₂CO₃(aq) + H₂O ⇌ H₃O⁺(aq) + HCO₃⁻(aq)
      HCO₃⁻(aq) + H₂O ⇌ H₃O⁺(aq) + CO₃²⁻(aq)

      CO₂ is always being produced in the cells as part of the normal metabolic process.
      It enters the blood, where it becomes part of the carbonate-bicarbonate buffer system.
      The blood travels to the lungs, where the dissolved CO₂ crosses the lung membranes and is exhaled in the breath.
      (7 votes)
  • leaf yellow style avatar for user saransh60
    Ok I got it when we add more H+ into the solution the equilibrium reaction is going to go to left resulting in the formation of more carbonic acid. More carbonic acid means that CO2 is less likely to react with water to form carbonic Acid, but what tells the CO2 to stop reacting with H20 and tell hey we already have enough Carbonic acid formed already .
    (4 votes)
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    • old spice man green style avatar for user Matt B
      This is a basic property of equilibrium. In chemistry, nothing is really "static" and in fact most reactions are "dynamic". What happens is that the forward reaction is initially much greater than the backward reaction but after a certain time, closer to equilibrium levels, the forward reaction becomes small enough (because there are fewer reactants) and it becomes equal to the backward reaction. Once the forward is equal to the backward reaction, it seems like things have stopped reacting.
      (2 votes)
  • primosaur sapling style avatar for user 0xB01b
    what is another example of a bicarbonate
    (3 votes)
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  • hopper cool style avatar for user 🤔 ᴄᴏᴅᴇᴅ ɢᴇɴɪᴜȿ 😎
    Are acidemia and alkalosis bad? For what Sal said they were, they didn't seem too bad...
    Also, what is a better description of them?
    (3 votes)
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  • blobby green style avatar for user annabel
    If I add a base, then the concentration of hydrogen ions will go down. Presumably they are being made into water. Where do the water molecules go? Doesn't this water affect the pH balance of the blood?
    (3 votes)
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    • old spice man green style avatar for user Matt B
      Nice question! If you add a base then the [H+] will go down and the pH will increase. This is the primary effect. Yes, more water will then start existing but the amount of water created compared to the concentration of acid/base particles is insignificant.

      Example: if you had a concentration of 10^-7 moles of H+ (pH 7) and suddely have 10^-8 moles (pH 8) as a consequence of adding a base, then you only increase the relative concentration from 1 to approx. 1.00000001.
      (3 votes)

Video transcript

- [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.