# Blood pH



## mikie (Nov 22, 2010)

We were talking about blood pH today (mostly in regards to changes from hypovolemia).  

Teacher: "the blood pH should be between 7.35-7.45"

My question:

Is there a difference in pH of venous vs arterial blood?

My rationale:
Venous pH would be higher because of the metabolic byproducts being exhausted by the tissues.  

Teacher wasn't sure of the "textbook" answer.  Searching Google yielded me nothing but medical articles that require an expensive subscription or not relevant.

Thanks

{paging the biochemistry people too}


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## usalsfyre (Nov 22, 2010)

Would your pH be higher or lower? Think about what cells are exhausting and research ABG vs VBG values.


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## socalmedic (Nov 22, 2010)

short answer: no

long answer: there may be a slight variations in blood pH post capillary exchange. however this will be very slight in a healthy person. this is why you get the pH from the ABG, it is standardized. in all reality the variations will not alter our treatment at all (keep in mind that the lab values may extend out to the thousands "7.438" or further), maby JP can weigh in on this? he most likely has access to more reputable sources than curious gurney slingers. 

remember though that most field providers are limited to non invasive measurements (capnography) to measure blood pH, in which 35-45mmHg is normal EtCo2


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## JPINFV (Nov 22, 2010)

Ahh, I thought I replied already, but I never hit the button...

D: None of the above because the buffer system is designed to maintain a pH between 7.35 and 7.45. If anything, though, the pH of venous blood should be lower due to an increase in CO2 pushing the buffer system towards the left, but then again, a buffer system is characterized by substances (sodium bicarb <->H2O and CO2) that absorb or release hydrogen depending on the amount of free hydrogen.


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## Sandog (Nov 23, 2010)

Also interesting to note albeit a lesser role than bicarb, the de-oxygenated hemoglobin (unbounded) proteins have an affinity for H+ protons.The excess protons bind to the hemoglobin thus increasing pH. The H+ unbinds when O2 attaches to the heme group.


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## mikie (Nov 23, 2010)

Why wouldn't the acids in the venous system (lactic, pyruvic) increase the pH?

JPINFV (and anyone else who can answer)- talk to me about this buffer system or throw a link please; i'm very interested.


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## Veneficus (Nov 23, 2010)

mikie said:


> Why wouldn't the acids in the venous system (lactic, pyruvic) increase the pH?
> 
> JPINFV (and anyone else who can answer)- talk to me about this buffer system or throw a link please; i'm very interested.



The more acidic, the lower the PH.

PH 1 is more acidic 
PH 9 is more basic.

Blood buffering is rather complex, probably too complex for a link (you know, in addition to the cellular specifics, hematology, there is a renal component, and compartment shifts and hormones.), but I can refer you to Guyton's Textbook of Medical Physiology.


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## mikie (Nov 23, 2010)

When I say in increase in pH I'm saying a rise in Hydrogen Ion concentration, not necessarily the #...


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## fma08 (Nov 23, 2010)

mikie said:


> Why wouldn't the acids in the venous system (lactic, pyruvic) increase the pH?
> 
> JPINFV (and anyone else who can answer)- talk to me about this buffer system or throw a link please; i'm very interested.



That goes back to math, (gross I know). pH is based on a negative logarithmic scale. pH=-log[H]. As Vene pointed out, the more H, the lower the number, the less H, the higher the number. Also keep in mind that it is a log scale, a jump from 6-7 is a very large jump in concentration.

Specifically, why an addition of acid would not increase or lower the pH (at least significantly) again is due to the buffer system in the  blood. It is there to keep the pH in the normal range so proteins and such don't start losing their function.

Buffers 101:
http://www.youtube.com/watch?v=g_ZK2ABUjvA&feature=related

The Body Version:
http://www.youtube.com/watch?v=HrUvft2d8Zo&feature=related
http://www.youtube.com/watch?v=-8tcJ8uBn-Q&feature=related


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## JPINFV (Nov 23, 2010)

mikie said:


> When I say in increase in pH I'm saying a rise in Hydrogen Ion concentration, not necessarily the #...



The problem is that's the absolute opposite of what a rise in pH means.


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## mikie (Nov 23, 2010)

JPINFV said:


> The problem is that's the absolute opposite of what a rise in pH means.



Acid ions concentration goes up=low pH #
Low Acid Ion concentration=High pH# 

Right?


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## fma08 (Nov 23, 2010)

mikie said:


> Acid ions concentration goes up=low pH #
> Low Acid Ion concentration=High pH#
> 
> Right?



Yes, but it may not be so cut and dry with the buffer systems in place.


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## JPINFV (Nov 23, 2010)

fma08 said:


> Yes, but it may not be so cut and dry with the buffer systems in place.



Kinda of, sorta. pH is all about the equilibrium between free hydrogen ions (technically H3O) and hydrogen bound to it's conjugate base. The more likely that a hydrogen ion will disassociate from its base, the lower it will drive the pH. With a buffer solution the amount of hydrogen ions that will associate or disassociate varies greatly over a small range of free hydrogen concentration (pH). Essentially, a buffer is like a hydrogen sponge. The hydrogen ions are still there, but if they're in the sponge, they aren't contributing to the pH.


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## fma08 (Nov 23, 2010)

JPINFV said:


> Kinda of, sorta. pH is all about the equilibrium between free hydrogen ions (technically H3O) and hydrogen bound to it's conjugate base. The more likely that a hydrogen ion will disassociate from its base, the lower it will drive the pH. With a buffer solution the amount of hydrogen ions that will associate or disassociate varies greatly over a small range of free hydrogen concentration (pH). Essentially, a buffer is like a hydrogen sponge. The hydrogen ions are still there, but if they're in the sponge, they aren't contributing to the pH.



That's what I was saying.

No buffer: increase in H+ (H3O) = decrease pH, decrease in H+ = increase in pH

With buffer: not so cut an dry, increase in H+ may or may not decrease the pH, any buffer system can be overcome if enough H+ is added beyond the saturation point of conjugate salt in the solution.


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## Sandog (Nov 23, 2010)

fma08 said:


> That's what I was saying.
> 
> No buffer: increase in H+ (H3O) = decrease pH, decrease in H+ = increase in pH
> 
> With buffer: not so cut an dry, increase in H+ may or may not decrease the pH, any buffer system can be overcome if enough H+ is added beyond the saturation point of conjugate salt in the solution.



Well sure, but hopefully acid-base homeostasis in the body prevents that situation from getting out of control. Otherwise we got a very sick person.


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## fma08 (Nov 23, 2010)

Sandog said:


> Well sure, but hopefully acid-base homeostasis in the body prevents that situation from getting out of control. Otherwise we got a very sick person.



We're not here for healthy people


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## silver (Nov 23, 2010)

mikie said:


> We were talking about blood pH today (mostly in regards to changes from hypovolemia).
> 
> Teacher: "the blood pH should be between 7.35-7.45"
> 
> ...




Also look up the Bohr effect to learn more about the wonders of the hemoglobin, and how H+ concentration changes and 2,3-BPG alter the affinity of hemoglobin for oxygen.

Great way to learn about allostery incase you are interested in biochem.


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## medichopeful (Nov 24, 2010)

mikie said:


> Why wouldn't the acids in the venous system (lactic, pyruvic) increase the pH?
> 
> JPINFV (and anyone else who can answer)- talk to me about this buffer system or throw a link please; i'm very interested.



I know there were some links, but here's a very quick and basic run-down of the bicarbonate-buffer system.

The normal equilibrium equation for normal blood is as follows:
(H20)+(CO2) <-> (H2CO3)+(H2O) <-> (HCO3-)+(H3O+)

If it works correctly, the CO2 equalizes through carbonic acid (H2CO3) into hydronium ions (H3O+) and bicarbonate ions (HCO3-).  When it works, this keeps the pH of the blood between 7.35 and 7.45.  When it doesn't work, the pH can change, as Le Chetelier's principle takes effect.

When it doesn't work, there are 4 basic things that can happen: respiratory acidosis, respiratory alkalosis, metabolic acidosis, and metabolic alkalosis.  

Respiratory acidosis:
This results when the patient has too much CO2 in their system, which can result from hypoventilation or something else.  This causes the equilibrium reaction to proceed like this:
(H20)+(CO2) -> (H2CO3)+(H2O) -> (HCO3-)+(H3O+)
In other words, more CO2 leads to an increase in the amount of carbonic acid which leads to an increase in the amount of bicarbonate ions and hydronium ions.  An increase in hydronium ions leads to a drop in the pH of the blood (provided the buffer system can't overcome it).

Respiratory alkalosis:
This is along the same lines as respiratory acidosis, except instead of an increase in CO2 there's a deficit of CO2.  This results in the following equilibrium reaction:
(H20)+(CO2) <- (H2CO3)+(H2O) <- (HCO3-)+(H3O+)
The reaction shifts to the left, to make up for the lack of CO2.  This reduces the amount of hydronium ions, which results in a higher (and thus more basic) pH.

Metabolic alkalosis:
This results when a condition causes the pH of the blood to rise, and thus become more basic.  This results in the following equilibrium equation:
(H20)+(CO2) -> (H2CO3)+(H2O) -> (HCO3-)+(H3O+)
Due to the drop of the pH on the right side of the equation, the equation shifts to the right to try to balance itself out.  If the buffer system can't overcome the change in alkalinity, it can effect breathing and other things.

Metabolic acidosis:
This is the final one.  When this occurs, such as in diabetic ketoacidosis, there is an excess of acid put into the bloodstream.  This causes the equation to shift to the left:
(H20)+(CO2) <- (H2CO3)+(H2O) <- (HCO3-)+(H3O+)
The increase in acidity of the blood causes the amount of CO2 to increase, which causes hyperventilation as the body tries to rid itself of the extra CO2.

I hope this helps!
Eric


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## mikie (Nov 24, 2010)

Thanks for all the replies!


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## Smash (Nov 25, 2010)

medichopeful said:


> I know there were some links, but here's a very quick and basic run-down of the bicarbonate-buffer system.
> 
> The normal equilibrium equation for normal blood is as follows:
> (H20)+(CO2) <-> (H2CO3)+(H2O) <-> (HCO3-)+(H3O+)



You seem to have picked up some extra fluid somewhere. (H2O) + (CO2) <> (H2CO3) <> (HCO3-) + (H+) would be the result of your equation, and this is commonly used to demonstrate the carbonic-acid-bicarbonate equilibrium.  Your equation needs an extra H2O on the left to make it correct and show the 2 equations, although to be honest I'm not sure that for the purposes of your average ambulance driver that it adds any useful info.


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## medichopeful (Nov 25, 2010)

Smash said:


> You seem to have picked up some extra fluid somewhere. (H2O) + (CO2) <> (H2CO3) <> (HCO3-) + (H+) would be the result of your equation, and this is commonly used to demonstrate the carbonic-acid-bicarbonate equilibrium.  Your equation needs an extra H2O on the left to make it correct and show the 2 equations, although to be honest I'm not sure that for the purposes of your average ambulance driver that it adds any useful info.



I've seen the equation written both ways, your way and mine.  The way I wrote it [(H20)+(CO2) <-> (H2CO3)+(H2O) <-> (HCO3-)+(H3O+)] shows that there is extra fluid in the body, shows the role of water, and is not balanced.  I see what you mean by saying that I would need an extra H2O on the left, but as this equation is not balanced (it just takes into account the excess fluid in the body), it's actually not necessary.  Your way is probably a little bit clearer, because it takes water out of the equation.  But the fact of the matter is, the water is still there in yours, you just don't see it in the equation.  I hope this makes sense :wacko:


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## MrBrown (Nov 25, 2010)

Oh bloody hell Brown knew finishing reading that basic chemistry book would be a good idea... where is Mrs Brown when you need her, after all she is the chemistry smartypants at Cassa de Brown.


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## JPINFV (Nov 25, 2010)

Smash said:


> You seem to have picked up some extra fluid somewhere. (H2O) + (CO2) <> (H2CO3) <> (HCO3-) + (H+) would be the result of your equation, and this is commonly used to demonstrate the carbonic-acid-bicarbonate equilibrium.  Your equation needs an extra H2O on the left to make it correct and show the 2 equations, although to be honest I'm not sure that for the purposes of your average ambulance driver that it adds any useful info.



H2O is in excess and H+ vs H3O (hydronium ion) is the same thing as you won't have free hydrogen ions floating around when water is present.


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## westcoastmedic (Nov 25, 2010)

Regardless of some of the incorrect information I found this topic very interesting and informative. Happy Turkey Day!!!


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## Smash (Nov 25, 2010)

Sorry, I'm probably being a touch pedantic. The first equation is used to show the effect of water on the acid-base and non acid-base equilibrium, but to be correct it should read (2H2O) + (CO2) <> ..... and so on. Otherwise it is mathematically incorrect (biology just being chemistry, chemistry being physics and physics being maths after all  )


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## Sandog (Nov 25, 2010)

JPINFV said:


> H2O is in excess and H+ vs H3O (hydronium ion) is the same thing as you won't have free hydrogen ions floating around when water is present.



Now ya done it, I'm gonna have Titration lab nightmares again. Blip, blip...


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## mikie (Nov 28, 2010)

*Paging Einstein!*

OK so as I'm reading through this and trying to decipher what you chem whiz'z said (though I did take my fair share of chemistry)...

This is occurring in the capillary beds? Venous system? Diffusing during respiration?  ----Ya those are loaded questions, but I'm just trying to understand how Arterial blood and Venous blood have similar pH's.  Yes it's obviously complex medicine beyond the SOP for medics, I'm just trying to learn B) and understand how this is possible (on a simpler scale). 


Thanks EVERYONE!!


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## Sandog (Nov 29, 2010)

Here is some good text on the subject.

http://www.madsci.com/manu/gas_acbs.htm

Also, here is a article on a Comparison of arterial and venous blood gas values.

http://www.rmj.org.pk/rmj_jan_jun_2008/original_articles/comparison_of_arteial_and_venous/pdf.pdf


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## Commonsavage (Nov 29, 2010)

*Great discussion, but....KISS*

I think the simple answer to a simple question was lost in the minutea of acid-base explanations, some of which are nothing more than complicated guesswork.
So, I will attempt the simple:
1) pH is a relative measurement of the ratio of hydrogen ions (H+) to hydroxide ions (OH-)
2) the pH scale is 1 to 14.  7 is neutral, in which the ratio of H+ to OH- is 1:1.
Anything below 7 is acid, higher in hydrogen ions.  Anything above 7 is basic (alkylotic), higher in hydroxide ions.
3) Venous blood is relatively more acid than arterial blood due to the way in which carbon dioxide is carried in the venous serum, in the form of carbonic acid.  The presence of lactic acid or pyruvate has very little to do with it.  It is the accumulation of CO2 in the form of carbonic acid that is the major and overriding factor.
What you need to know is that, metabolic and/or respiratory issues that take the human organism outside of the narrow limits of 'normal' pH (7.35 - 7.45) can be critical and can be corrected...within limits.  Outside of the correctable limits (6.5 - 7.65), recovery is doubtful.  Acidosis is much more prevalent than alkylosis.  As one's condition deteriorates, with decreased respiratory status and increased tissue catabolism, one becomes acidic.

That, my friend, is the VERY basic bare bones of pH blood chemistry.  I am really sorry that your instructor is so ill educated.  Unfortunately, at least in my observations of EMT/Paramedic certification/education programs in this country, we have people placed as instructors in courses about which they know only enough to get in trouble, generate myths, and pass on bad information.  By the way, that's pretty much the status of many nursing instructors as well.


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## mikie (Nov 30, 2010)

thanks for the links; Sandog

and i like your summary, Commonsavage!


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