Carbonic acid bicarbonate buffer system
By: Karim Ali and Adam Abdelghaffar
19th of May, 2025
Ap Chemistry
Introduction
The system we chose is a carbonic acid bicarbonate buffer system. This system is an example of
an Acid-Base equilibrium system, this is because it occurs
naturally in the human body. It involves a reaction between
Carbon dioxide ( CO2), water (H2O), carbonic acid
(H2CO3), bicarbonate (HCO3-), and hydrogen ions (H+).
The goal of this Acid-Base equilibrium reaction is to
stabilize the blood pH in the body, this is done by the
Bicarbonate (HCO3-) buffering the excess acid in the body, which is extremely crucial for overall
metabolism and enzyme function.
Steps for how the reaction works
1. Carbon dioxide (CO2) is produced: This is when the carbon (C) and the oxygen (O2) in
the Carbon dioxide (CO2) atoms come from the cellular respiration ( breakdown of
glucose). The CO2 that is formed is then released into the bloodstream as a waste product,
which then can buildup and lead to a more acidic pH if not removed or buffered.
C6H12O6 +6O2 → 6CO2 + 6H2O + ATP
2. Formation of Carbonic Acid (H2CO3): The CO2 that is produced as a waste product, is
then diffused into the blood of the body and then enters the red blood cells. The CO2 that
was produced, enters the red blood cells and reacts with existing H2O atoms that are
already present in the blood plasma and red blood cells, which then produces the
Carbonic acid.
CO2 + H2O → H2CO3
3. Carbonic acid (H2CO3) dissociates: This is when the unstable carbonic acid atoms break
apart into different atoms, into hydrogen ions (H+) and bicarbonate ions (HCO3-). The
hydrogen ions affect the blood pH by making it more acidic, and the bicarbonate ions act
as a buffer to neutralize the blood pH to its normal level. This dissociation happens
naturally in blood when there is more CO2 present, CO2 is present when someone is
exercising or holding their breath.
H2CO3⇌H+ + HCO34. Buffering excess H+ (acidic conditions): This is when the hydrogen ions (H+) and the
HCO3- recombine to create H2CO3, which then breaks into two different compounds: CO2
and H2O. Then the body neutralizes the excess H+, which helps remove the acidity from
the blood in order to restore the normal pH range for human blood, this helps prevent
dangerous pH blood levels (acidosis).
H+ + HCO3- → CO2 + H2O
5. Buffering low H+ (basic conditions): This is when the CO2 and H2O that come from the
cellular respiration form carbonic acid, which releases H+ into the blood of the body, in
order to raise the blood pH, this is in order to prevent the blood pH being too basic
(alkalosis)
6. Lungs and kidneys help regulate levels: Here the lungs control the CO2 levels by
adjusting the breathing rate that someone is breathing with, the faster someone breathes
the more CO2 is removed, which results in less acid being built up in the blood and
prevents acidosis. The slower someone breathes the less CO2 is removed, so the
individual retains the amount of CO2 in their body, which results in the blood becoming
too acidic (acidosis). The kidneys help remove H+ ions in an individual's urine if their
blood is too acidic, this is in order to keep blood pH levels in the normal range. The
kidneys also reabsorb or remove HCO3- to adjust the buffering power of the system, so an
individual's blood is neither too acidic or basic, preventing acidosis and alkalosis.
Literature Review
The Bicarbonate buffer system is an essential process that helps to regulate the acid-base balance
that is in the blood and in other fluids in our body. It balances the levels of carbonic acid, carbon
dioxide, and the bicarbonate ions in the human body to maintain a stable and well physiological
task. Equilibrium in the system is maintained through a reversible reaction that occurs between
carbonic acid (H₂SO₃) and bicarbonate (HCO3-). The system shifts based on the pH changes of
the carbonic acid and bicarbonate. When the blood becomes too acidic, the bicarbonate reacts
with excess hydrogen ions in order to form carbonic acid, which results in reducing the acidity of
the blood. But, if the blood is too basic, carbonic acid tends to dissociate to release hydrogen
ions, which increases the acidity. Carbon dioxide levels in the body affect the concentration of
carbonic acid. An increase in carbon dioxide results in more carbonic acid, thus helping in
lowering the pH. While a decrease in carbon dioxide leads to less carbonic acid, increasing the
pH. There are numerous factors that can influence the position of equilibrium in the buffer
system. One of the factors is the concentration of the reactants and products, where changes in
levels of both carbonic acid and bicarbonate can shift the equilibrium, an increase in hydrogen
ions will make the reaction favored towards the carbonic acid and a decrease will move the
reaction towards bicarbonate. Another factor is the partial pressure of carbon dioxide. If there are
higher levels of carbon dioxide, the amount of carbonic acid would increase, leading to a
leftward shift in the equilibrium and a decrease in the pH. A decrease in carbon dioxide levels
would shift the equilibrium to the right, highering the pH. An increase in temperature would
raise the formation of bicarbonate, affecting the buffer capacity.
Hypothesis development
As temperature, pressure and concentration are factors that affect the equilibrium equation CO2 +
H2O ⇌H2CO3 , an increase or decrease in the quantity will affect the positions of the chemicals
in the equation. An increase in temperature will move the reaction towards the reactants in order
to absorb the heat, which results in a decrease in the concentration of H2CO3 , while a decrease in
temperature will result in the opposite as the products will be favored. Moving on to
concentration. An increase in the reactants’ concentration will result in the equilibrium shifting
to the right so the reaction balances. But, an increase in the products’ concentration will cause
the reaction to shift to the left, and will favor the formation of CO2 and H2O. The equilibrium
will shift towards the side where the concentration is less to balance the equilibrium. A change in
pressure mainly affects the gases in a reaction. If the pressure increases, the equilibrium will shift
towards the side that has less gas molecules. Regarding this equation, the reaction will move
towards the right as there are no gas molecules. If the pressure decreases, equilibrium will move
to the left, resulting in the favoring of the reactants.
Experimental design
Goal of the experiment:
Investigate how changing the concentration of CO2 affects the equilibrium position of the
carbonic acid-bicarbonate system.
Chemical reaction:
CO2 (aq) + H2O (l) ⇌H2CO3 (aq) ⇌H+ (aq) + HCO3- (aq)
In this experiment the independent variable is the concentration of CO2, this is because it is the
one changing and the dependent variable is the pH of the solution, this is because it is dependent
on the concentration of the CO2. The controlled variables are the temperature, volume of water,
time for CO2 to bubble, and the type of container.
Materials:
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Distilled water ( at least 1 litre)
Carbon dioxide (CO₂) gas
pH indicator solution
5 Erlenmeyer flasks or beakers (100–250 mL each)
Gas syringe or CO₂ delivery setup
Digital pH meter
Graduated cylinders (100 mL, for measuring water)
Thermometer
Water bath or hot plate
Stopwatch or timer
Rubber tubing and corks
Tripod stand and clamp
Experimental setup:
1. Prepare 5 identical beakers with 100 mL distilled water
2. Fill the syringe with the desired volume of CO2 gas (50 ml, 100ml, 150ml, 200ml, and
250ml)
3. Then bubble the different volumes of the CO2 gas each for 5 minutes
4. After the bubbling is finished, measure the pH
5. Repeat for each volume and compare the pH
6. Repeat the experiment 3 times for reliable results, then use the average pH from the 3
trials
Record your results in this table
Hypothesis
We expect that increasing the CO2 will shift the equilibrium to the right, this is because
according to le chatelier's principle, increasing one of the reactants will cause the reaction to
adjust itself by producing more reactants, meaning that the reaction will move to the right and go
to completion since the reactants and becoming less and the products are increasing. This will
cause the formation of more hydrogen ions (H+), which then lead to the pH to decrease. This led
to us concluding that as the CO2 concentration increases the pH will decrease, meaning that the
CO2 concentration and pH are inversely proportional. This will all happen because more CO2
dissolves in the water, which will cause more formation of the weak acid carbonic acid (H2CO3),
then the carbonic acid breaks down into H+ and HCO3-, causing the H+ concentration to increase,
which is the reason for why the pH is decreasing. This tells us that the increase of CO2
concentration will increase the H+ concentration, leading to the decrease in pH causing the blood
to be more acidic. In order to get rid of the H+ in the blood that is causing it to be acidic, you will
need to manipulate the carbonic acid bicarbonate buffer system, so you are able to create a buffer
system that resists the H+ ions and helps lead to a more neutral blood pH. We will talk about this
more in our real world application part of the research paper.
This graph shows that our hypothesis was right, that as CO2 volume increased the pH decreased
making them inversely proportional.
Analysis of real world application
1. Athletes and baking soda (sodium bicarbonate)
a. Athletes that require high intensity effort for a long period of time, usually 1-7
minutes of performance, like swimming, boxing, and sprinting, sometimes
consume baking soda (NaHCO3) 90 minutes before they perform. The athlete can
only consume up to 0.3 grams per kg of their body weight, because too much
sodium bicarbonate can cause distress and ultimately affect their performance
rather than elevating their performance.
b. This is because as they produce their maximum amount of effort, the lactic acid in
the blood increases greatly, causing the blood pH to drop significantly and
causing acidosis. This can cause their performance to drop during the rest of the
performance, which can cause them to not get the results they desired from their
performance.
c. To avoid this the sodium bicarbonate they consumed 90 minutes before their
performance acts as a buffer in the reaction. This helps neutralize the excess H+
that was produced from the formation of lactic acid and helps maintain the
optimal blood pH, which then delays the fatigue that causes their performance to
drop. This ultimately helps athletes to increase their performance yielding greater
results from their performance, by just simply consuming a small amount of
baking soda (sodium bicarbonate).
H+ +HCO3- → H2CO3 → CO2 + H2
2. Medical application - Blood pH regulation
a. In hospitals, some patients that come in with respiratory distress often can not
exhale the excess CO2 in their body properly, which causes the excess CO2 to
build up in their blood, which not only shifts the equilibrium to the right
increasing the H+, but also causes their blood pH to decrease causing acidosis.
b. Doctors are able to prevent this by monitoring their arterial blood pH and if it
becomes too acidic, they use ventilators in order to help the patient to exhale the
excess CO2 in order to regulate the blood pH in their body. They also sometimes
add bicarbonate intravenously in extreme cases so they are able to neutralize the
acidity.
c. This is in order to keep the pH from 7.35-7.45 which is the safe zone of blood pH,
avoiding acidosis or alkalosis.
Overall, in both applications the bicarbonate buffer system acts as a way to neutralize the excess
H+ ions, so the blood is restored to its normal range in order to prevent acidosis or alkalosis. For
athletes, this helps boost their endurance and performance by buffering the lactic acid. For
patients in the hospital, this helps support their breathing and stabilizes patients in critical
conditions. This all goes to show the importance of the carbonic acid bicarbonate buffer system,
it not only helps people perform better at sports, but also help save lives: showing how critical it
is to keep the blood pH stable and in the normal range.
Discussion
Importance of understanding equilibrium
When a condition changes, like the temperature, addition of a reactant, addition of a product, etc.
The equilibrium shifts to oppose that change and help reestablish equilibrium. Adding a reactant
will cause the system to shift towards the product and get rid of the reactants. Removing a
product will do a similar thing, it will cause the reaction to shift towards the products in order to
replace the products that were removed, so equilibrium is reestablished. Equilibrium is important
to maintain stability in living systems, in this case in humans. This equilibrium equation is
important, so we are able to regulate blood pH, ensuring enzymes, oxygen transport, and
metabolic processes function correctly. Even the slightest deviation in pH outside of 7.35-7.45
can cause acidosis and alkalosis, which can cause distress for an individual. It is also important
because it helps us predict what will happen in the reaction, this is done by using Le Chatelier’s
principle, which we will talk about more later in the paper. Lastly, equilibrium helps support
industrial and environmental processes. In industrial processes, it helps control equilibrium
conditions in order to help maximize yields like ammonia production in the haber process, and in
environmental science, equilibrium principles help explain how the ocean becomes acidic, where
the CO2 lowers the seawaters pH, exactly like the human body.
Prediction of the outcome using Le Chatelier’s principle
● When a condition changes, as we stated before in the importance of understanding
equilibrium, the equilibrium opposes that change to reestablish equilibrium. For example,
adding CO2 will cause the reaction to shift right, producing more H+, causing the pH to
drop. Removing CO2 does the opposite, causes the reaction to shift left, producing less
H+ and the pH rises. Adding H+ combines with HCO3- to remove H+ and neutralize the
pH in the body. Adding OH- does the opposite of adding H+, H+ neutralizes with the OH-,
causing the equilibrium to shift in order to produce more H+, causing the pH to drop.
● Understanding all of this is extremely important, this is because it helps you predict what
will happen when a certain condition happens, not only that but it makes you able to
manipulate the reaction to benefit you instead of affecting you: like we explained in the
two real world application and how they manipulated the reaction to benefit athletes and
patients at hospitals. It also helps you understand the chemical balance in your own body,
and what you can do to balance the blood pH in your body, in order to keep your blood
pH in the safe zone.
Conclusion
After researching and experimenting with the carbonic acid bicarbonate system, we found out
that the equilibrium principles for this reaction especially are critical for maintaining life
sustaining blood pH with the rage being 7.35-7.45. The findings of our experiment confirmed for
us that the increase in CO2 concentration decreases the pH, which is extremely bad for the
human body. This is because the decrease in blood pH can cause acidosis, which can cause
distress or even life threatening dangers. This helps support the idea of Le chatelier's principle,
that with any change in the standard conditions, the reaction will shift itself in order to
reestablish equilibrium. The understanding of Le chatelier's principle and equilibrium, led to
doctors, engineers, scientists, and athletes to manipulate the reaction in order to help them with
their performance or task that they have been assigned with in order to yield the best results
possible. This was explained during our research, using two real life applications. For athletes,
they manipulate this reaction in order to delay fatigue and improve their performance, and for
doctors they manipulate the reaction to help patients with respiratory issues. This leads us to
conclude that our research emphasizes that chemical equilibrium is not just a concept in
chemistry, but it is something extremely important and when understood, can be used to our
advantage so we are able to perform better at our jobs or help people in distress. Displaying the
importance of the carbonic acid bicarbonate buffer system in our life.
Citations
1. Khan Academy. Buffers in the Human Body. Khan Academy,
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LibreTexts, 2 Oct. 2013,
https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_General_Chemistr
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m. Accessed 21 May 2025.
3. Carr, A. J., et al. “Sodium Bicarbonate Supplementation and Exercise Performance.”
Sports Medicine, vol. 41, no. 9, 2011, pp. 733–750. PubMed,
https://pubmed.ncbi.nlm.nih.gov/29562565/. Accessed 21 May 2025.
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National Center for Biotechnology Information, 2009,
https://www.ncbi.nlm.nih.gov/books/NBK54104/. Accessed 21 May 2025.
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