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Been struggling with this for the past few days even after reading half of the acid base tutorial here, if someone could help me that'd be great.

What I don't understand is how HCO3- is supposed to buffer the blood. I know that HCO3- is a base and can accept H+, but the way that the HCO3- came to fruition is by donating a H+ to the surrounding water, so there's a net zero change in H+ thus no buffering action took place. Let's say a HCO3- accepts a H+, now the CO2 produced by H2CO3 → CO2 + H2O leaves the blood because the PvenousCO2 gets higher than the CO2 pressure in the air in the lungs, and you exhale. Now you're in the exact same situation as before the CO2 got attached to H2O. Still the text calls it as an buffer, with a very large buffering capacity as well. How does this work?

To me it seems that buffering blood using HCO3- makes as much sense as using HCl, Cl - can accept H+, right? Nevermind that HCl is a very strong acid.. Surely no one would think of using HCl to buffer blood.

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  • $\begingroup$ I just have one question for you. If the body could regulate the amount of HCO3-, in effect generating new buffer, would that help you to understand how it is the major buffer that it is? $\endgroup$ – anongoodnurse Dec 27 '14 at 18:25
  • $\begingroup$ maybe, not sure if I understand it now $\endgroup$ – Plumpie Dec 31 '14 at 13:56
  • $\begingroup$ Wish you had answered me days ago. $\endgroup$ – anongoodnurse Dec 31 '14 at 20:15
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What do you mean by "there's a net zero change in $\ce{H+}$ thus no buffering action took place"? The property of a buffer system IS to prevent drastic change in pH.

Buffer capacity, a measure of how much a buffer system can sustain a change in $\ce{[H+]}$ (aka. pH), (mostly) depends on two parameters -- the concentration of that buffer, and how close the desired pH is to its pKa. $\ce{H2CO3}$ has pKa's of 6.4 and 10.3. Given that the pH of blood is about 7.4 and better to stay that way, the base/acid pair, $\ce{[HCO3_^{-}]/[H2CO3]}$ (pKa=6.4) (plus a whole bunch of various proteins such as albumins in blood, all of which has its own pKa) does provide a good buffer capacity around that pH.

True, $\ce{CO2}$ is constantly removed through respiration, but it is also constantly generated by cellular respiration, making $\ce{H2CO3<=>CO2 + H2O}$ a "volatile equilibrium" and keeping both $\ce{HCO3_^{-}}$ and $\ce{H2CO3}$ in high enough concentration to sustain its buffer capacity.

HCl has a pKa of about -7 so it would make a good buffer around pH of -7. In an aqueous environment, all HCl molecules will 100% dissociate, and no $\ce{H+}$ will ever bind to $\ce{Cl-}$ because water itself will first "hijack" the proton as $\ce{H3O+}$ (pKa=-1.74) before $\ce{Cl-}$ by about five order of magnitude.

---EDIT---

You may be thinking that when cells release $\ce{CO2}$ to the blood, $\ce{H2CO3}$ would be the first product, and in order for $\ce{H2CO3}$'s conjugate base to work as a base to absorb any extra $\ce{H+}$, first it must release a $\ce{H+}$, resulting in no net change in free $\ce{H+}$ and hence no capacity in balancing extra or insufficient $\ce{H+}$ to maintain the pH. Indeed, if $\ce{H2CO3}$ is the ONLY chemical species in the blood when it would not be a buffer. A buffer requires at least two species -- a potential proton donor and a potential proton acceptor, both of which in high and comparable concentration. pH is a measure of external free $\ce{H+}$, a result of final proton balance of all the acid and base in the blood, dictated by the Henderson–Hasselbalch equation. Thus when $\ce{H2CO3}$ is released in the blood, the blood's basic pH will causes most of $\ce{H2CO3}$ to become $\ce{HCO3_^{-}}$, releasing free proton. Again IF that's all the $\ce{HCO3_^{-}}$ in the blood after being made from $\ce{H2CO3}$, the blood's pH will drop, but there is already abundant amount of $\ce{HCO3_^{-}}$ in the blood, helping the system to reach new equilibrium. Of course, the buffer capacity of blood is not infinite. If there is so much metabolic acid being produced that it shifts the ratio of $\ce{[HCO3_^{-}]/[H2CO3]}$ toward excessive acid, acidosis would occur.

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  • $\begingroup$ I'd hate to see what it takes to get water at pH of -7. $\endgroup$ – user137 Dec 29 '14 at 7:30
  • $\begingroup$ @user137 it need not be water.. It can be any solvent. Base/acid/salts etc are defined with respect to a solvent (which is water in most cases). If H₂SO₄ is the solvent then K₂SO₄ will be a base, not a salt. You can also imagine what if ammonia is the solvent. $\endgroup$ – WYSIWYG Dec 29 '14 at 8:58
  • $\begingroup$ What do you mean by "there's a net zero change in H+ thus no buffering action took place"? I mean that the H+ absorption is nice and all, but it's only made possible by producing a H+ at an earlier step. It's like saying glycolysis produces 4 ATP - it does of course but there's 2 ATP needed for glycolysis, so the net gain in ATP is only 2. I accepted your answer btw, thanks for the help (and commenters) but I'm not sure if I understand it any better now. $\endgroup$ – Plumpie Dec 31 '14 at 14:10
  • $\begingroup$ I think I have better understood your difficulty. You may be thinking that when cells release CO2 to the blood, H2CO3 would be the first product, and in order for H2CO3's conjugate base to work as a base to absorb any extra H+, first it must release a H+, resulting in no change in free H+ and hence no capacity in balancing extra or insufficient H+ to main the pH. Please see Edit. $\endgroup$ – SYK Dec 31 '14 at 17:16
  • $\begingroup$ I don't personally understand how you can discuss this buffering system without discussing the role of the kidneys. $\endgroup$ – anongoodnurse Dec 31 '14 at 20:13

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